METHODS    IN    CHEMICAL 
ANALYSIS 


ORIGINATED   OR   DEVELOPED 

IN    THE   KENT    CHEMICAL   LABORATORY 

OF   YALE   UNIVERSITY 


COMPILED   BY 

FRANK   AUSTIN    GOOCH 

PROFESSOR  OF  CHEMISTRY  AND  DIRECTOR   OF   THE   KENT  CHEMICAL 
LABORATORY  IN  YALE  UNIVERSITY 


FIRST  EDITION 


NEW   YORK 

JOHN    WILEY   &    SONS 
LONDON:   CHAPMAN   &   HALL,   LIMITED 


COPYRIGHT,  1912, 

BY 
FRANK  AUSTIN  GOOCH 


Stanhope  flfress 

F.    H.  GILSON  COMPANY 
BOSTON,  U.S.A. 


PREFATORY    NOTE 


THE  object  of  this  volume  is  to  present  concisely  the 
principal  results  reached  by  workers  in  the  Kent  Chemical 
Laboratory  of  Yale  University  in  the  investigation  and  develop- 
ment  of  methods  in  chemical  analysis.  In  the  account  of 
processes,  modified  or  original,  only  proved  procedure  and 
immediately  related  experimental  data  are,  as  a  rule,  given. 
For  further  details  in  respect  to  the  elaboration  of  processes, 
the  discussion  of  difficulties,  and  the  experimental  illustration 
of  the  effects  of  varying  the  prescribed  procedure,  references 
are  given  to  the  original  sources  from  which  this  summary  has 
been  compiled.  To  his  colleagues,  Professors  Philip  E.  Browning, 
R.  G.  Van  Name,  and  W.  A.  Drushel,  the  compiler  is  much, 
indebted  for  valuable  criticism  of  the  proof  sheets. 


575454 


CONTENTS. 


CHAPTER   I. 
APPLIANCES  AND  GENERAL  PROCEDURE. 

Mechanical  Processes.  —  The  determination  of  products  gaseous  at  ordinary 
temperatures  by  loss  of  weight,  i.  The  distillation  and  condensation  of  vola- 
tile products,  3.  The  distillation  and  absorption  of  volatile  products,  4;  ap- 
paratus with  ground  joints,  4;  apparatus  with  sealed  joints,  5.  The  removal 
of  volatile  products  from  material  to  be  reserved  for  treatment.  The  preven- 
tion of  mechanical  loss  from  solution  in  reactions  evolving  gaseous  products,  6. 
The  transfer  of  liquids  and  gases  under  pressure,  7.  A  convenient  form  of 
rotary  shaker,  9.  The  purification  of  precipitates  by  solution  and  reprecipita- 
tion,  10.  Electrolytic  processes.  —  The  rotating  cathode,  n.  The  filtering 
crucible,  13;  electrolysis  and  subsequent  filtration,  13;  electrolysis  with  filtra- 
tion, 16;  electrolysis  with  continuous  filtration,  17.  The  fixation  of  chlorine 
on  the  silver  anode,  20;  hydrochloric  acid,  silver  anode,  and  platinum  cathode, 
20;  sodium  chloride,  silver  anode,  and  mercury  cathode,  22.  lodometric 
processes.  —  The  standardization  of  iodine  solutions  by  the  action  of  metallic 
silver,  27.  Arsenic  trioxide  as  an  iodometric  standard,  29.  The  starch 
indicator  for  free  iodine,  29.  Standard  tartar  emetic,  38.  Processes  of 
oxidation.  —  Arsenic  trioxide  as  a  standard,  41 ;  standardization  without 
iodine,  41;  standardization  with  the  aid  of  iodine,  42.  The  gravimetric 
standardization  of  permanganate,  42.  The  loss  of  oxygen  in  oxidation  by 
permanganate,  42;  concentration  of  acid,  42;  hydrochloric  acid  with  ferrous 
salts,  48;  hydrochloric  acid  with  oxalic  acid,  50;  effect  of  other  chlorides,  52. 
Acidimetry  and  Alkalimetry.  —  The  use  of  succinic  acid  as  a  standard,  54. 
Organic  acids  and  acid  anhydrides  as  standards,  56.  The  use  of  the  iodide- 
iodate  mixture  and  the  estimation  of  iodine  evolved,  59;  determination  of  free 
acids,  59;  determination  of  alkali  hydroxides  and  carbonates,  60;  determina- 
tion of  acids  liberated  in  hydrolysis,  61.  The  use  of  the  bromide-bromate 
mixture  and  estimation  of  the  bromine  evolved,  70.  Reaction  of  iodine  with 
alkali  hydroxides,  70. 

CHAPTER   II. 
THE  ALKALI  METALS. 

Sodium.  —  The  detection  of  sodium,  74;  the  estimation  of  sodium  as  the 
pyrosulphate,  79.  Potassium.  —  The  spectroscopic  detection  and  determina- 
tion of  potassium,  80;  detection  of  potassium,  80;  determination  of  potas- 
sium, 83.  The  separation  and  determination  of  potassium  as  the  perchlorate, 


vi  CONTENTS 

88.  The  estimation  of  potassium  as  the  pyrosulphate,  92.  The  volumetric 
estimation  of  potassium  as  the  cobalti-nitrite,  93;  potassium  in  the  pure  salt, 
94;  potassium  in  fertilizers,  96;  potassium  in  soils,  97;  potassium  in  animal 
fluids,  urine,  blood,  lymph,  milk,  98.  Rubidium  and  Caesium.  —  The  spectro- 
scopic  determination  of  rubidium,  102.  Estimation  of  caesium  and  rubidium 
as  the  acid  sulphates,  106. 

CHAPTER   III. 
COPPER;  SILVER;  GOLD. 

Copper.  —  The  gravimetric  determination  of  copper  as  the  sulphocyanate, 
IO8;  separation  of  copper  from  bismuth,  antimony,  tin,  and  arsenic,  112. 
The  determination  of  copper  as  cuprous  iodide  and  separation  from  cadmium, 
114.  The  electrolytic  determination  of  copper,  116.  The  iodometric  estima- 
tion of  copper,  1 1 8.  The  determination  of  copper  by  titration  of  the  pre- 
cipitated oxalate  with  potassium  permanganate,  123;  precipitation  of  copper 
oxalate,  125;  solubility  of  copper  oxalate,  125;  prevention  of  supersaturation, 
129;  precipitation  in  presence  of  acetic  acid,  130;  separations  by  Peters'  pro- 
cedure, 131;  separations  by  the  method  of  desiccation,  132;  separations  in 
presence  of  acetic  acid,  134;  determination  of  copper  associated  with  lead,  135. 
Silver.  —  The  gravimetric  determination  of  silver  as  the  chromate,  136.  The 
electrolytic  determination  of  silver,  138.  The  iodometric  estimation  of  silver, 
based  upon  the  use  of  potassium  chromate  as  a  precipitant,  140.  The  iodo- 
metric determination  of  silver,  based  upon  the  reducing  action  of  potassium 
arsenite,  143.  Gold.  —  The  electrolytic  determination  of  gold,  145.  The 
iodometric  estimation  of  small  amounts  of  gold,  146.  The  colorimetric  deter- 
mination of  small  amounts  of  gold,  150. 


CHAPTER   IV. 
BERYLLIUM;  MAGNESIUM;  CALCIUM;  STRONTIUM;  BARIUM. 

Beryllium.  —  Ammonium  beryllium  phosphate,  153.  The  conversion  of 
beryllium  chloride  to  beryllium  oxide,  153.  The  separation  of  beryllium 
chloride  from  ferric  oxide,  154.  Magnesium.  —  The  determination  of  mag- 
nesium by  precipitation  and  ignition  of  ammonium-magnesium  carbonate,  154. 
The  determination  of  magnesium  as  the  pyrophosphate,  156.  The  arsenate 
process  for  the  separation  of  magnesium  and  the  alkalies,  158.  Calcium, 
Strontium,  Barium.  —  The  detection  of  barium  and  strontium  associated  with 
calcium,  and  lead,  160.  The  separation  of  barium,  strontium,  and  calcium  by 
the  action  of  amyl  alcohol  on  the  nitrates,  162;  detection  of  strontium  and 
calcium,  163;  separation  and  estimation  of  strontium  and  calcium,  164;  esti- 
mation of  barium  and  calcium,  166;  estimation  of  barium  and  strontium 
together,  and  of  calcium,  166.  The  separation  of  barium  and  strontium  by 
the  action  of  amyl  alcohol  on  the  bromides,  167.  The  estimation  of  barium 
as  the  sulphate,  168;  in  presence  of  hydrochloric  acid,  168;  in  presence  of 
nitric  acid  and  aqua  regia,  170;  purification  of  precipitated  barium  sulphate, 


CONTENTS  Vii 

172.  The  estimation  of  barium  as  the  chloride,  174;  precipitation  by  ether- 
hydrochloric  acid  mixture,  174;  precipitation  by  acetyl  chloride  in  acetone, 
175;  separation  from  calcium  and  magnesium,  177.  The  precipitation  of 
barium  bromide  by  ether-hydrobromic  acid  mixture,  179.  The  estimation  of 
calcium,  strontium,  and  barium  precipitated  as  oxalates,  180;  gravimetric 
determination  of  strontium  and  barium,  180;  titration  of  oxalates  by  per- 
manganate, 181. 

CHAPTER  V. 
ZINC;  CADMIUM;  MERCURY. 

Zinc.  —  The  estimation  of  zinc  as  the  pyrophosphate,  185.  The  conversion 
of  zinc  chloride  to  zinc  oxide,  186.  The  electrolytic  determination  of  zinc,  186. 
The  estimation  of  zinc  by  precipitation  as  the  oxalate  and  titration  with  potas- 
sium permanganate,  187.  Cadmium. — The  estimation  of  cadmium  as  the 
oxide,  1 88;  precipitation  as  carbonate,  188;  precipitation  as  hydroxide,  189. 
The  estimation  as  the  pyrophosphate,  190.  The  electrolytic  determination  of 
cadmium,  191;  deposition  from  the  sulphuric  acid  solution,  191;  deposition 
from  solutions  containing  acetates,  192;  deposition  from  solutions  containing 
cyanides,  193;  deposition  from  solutions  containing  pyrophosphates  or  ortho- 
phosphates,  194.  Mercury.  —  The  gravimetric  determination  of  mercury  as 
mercurous  oxalate,  195.  The  determination  of  mercury  by  titration  with 
sodium  thiosulphate,  196.  The  estimation  of  mercury  by  precipitation  as 
mercurous  oxalate  and  titration  of  the  excess  of  precipitant  with  permanganate, 
197.  The  titration  of  mercurous  salts  with  potassium  permanganate,  198. 

CHAPTER  VI. 
BORON;  ALUMINIUM;  LANTHANUM;  THALLIUM. 

Boron.  —  The  gravimetric  determination  of  boric  acid,  201 ;  the  use  of 
calcium  oxide  as  a  retainer,  201 ;  the  use  of  sodium  tungstate  as  a  retainer,  204. 
The  acidimetric  estimation  of  boric  acid,  205;  neutralization  of  stronger  acids, 
206;  strengthening  of  boric  acid  by  mannite,  208.  The  iodometric  deter- 
mination of  boric  acid,  210.  Aluminium.  —  The  determination  of  aluminium 
by  precipitation  with  ether-hydrochloric  acid/ 214;  separation  of  alumin- 
ium from  iron,  214;  determination  of  aluminium  and  beryllium,  216;  deter- 
mination of  aluminium  and  zinc,  216;  determination  of  aluminium  and  copper, 
217;  separation  of  aluminium  from  mercury  and  bismuth,  217.  Lanthanum. — 
The  estimation  of  lanthanum  precipitated  as  the  oxalate,  218.  Thallium.  — 
The  determination  of  thallium  as  the  acid  sulphate  and  as  the  neutral  sulphate, 
219.  The  gravimetric  estimation  of  thallium  precipitated  as  thallic  hydroxide 
by  potassium  ferricyanide  and  potassium  hydroxide,  220.  The  gravimetric 
estimation  of  thallium  as  the  chromate,  221.  The  iodometric  estimation  of 
thallium  by  precipitation  with  potassium  dichromate  and  determination  of 
the  excess  of  the  precipitant,  222.  The  estimation  of  thallium  by  the  action 
of  potassium  ferricyanide  in  alkaline  solution  and  of  potassium  permanganate 
in  acid  solution  upon  the  ferrocyanide  produced,  223. 


Viii  CONTENTS 

CHAPTER  VII. 
CARBON;  SILICON;  TITANIUM;  ZIRCONIUM;  CERIUM;  TIN;  LEAD. 

Carbon.  —  The  determination  of  carbon  dioxide  in  carbonates  by  loss,  225 ; 
expulsion  of  carbon  dioxide  by  the  action  of  acid,  225;  expulsion  of  carbon 
dioxide  by  ignition,  226.  The  precipitation  and  gravimetric  determination  of 
carbon  dioxide,  228.  The  iodometric  determination  of  carbon  dioxide,  231; 
carbon  dioxide  in  carbonates,  232.  The  combustion  of  organic  substances  in 
the  wet  way,  234;  carbon  content  by  the  permanganate  process,  234;  carbon 
content  by  oxidation  with  chromic  acid,  236;  carbon  dioxide  evolved  and 
oxygen  used,  239.  Silicon.  —  The  detection  of  silicon  in  silicates  and  fluo- 
silicates,  241.  Titanium. — The  determination  of  titanic  acid  by  reduction 
and  titration  with  potassium  permanganate,  242.  Zirconium.  — The  separa- 
tion of  zirconium  from  iron  by  volatilization  of  the  latter  in  hydrogen  chloride, 
244.  Cerium.  —  The  separation  of  cerium  from  other  cerium  earths  by  the 
action  of  bromine  upon  the  mixed  hydroxides  in  presence  of  an  alkali  hydrox- 
ide, 244.  The  iodometric  estimation  of  cerium,  246;  digestion  process,  246; 
distillation  process,  247.  The  estimation  of  cerium  oxalate  by  potassium  per- 
manganate, 248.  The  estimation  of  cerium  in  presence  of  other  rare  earths 
by  the  action  of  potassium  ferricyanide  in  alkaline  solution  and  potassium  per- 
manganate in  acid  solution,  249.  Tin.  —  The  electrolytic  determination  of 
tin,  251.  Lead.  —  The  detection  of  lead,  252.  The  electrolytic  determination 
of  lead  as  the  dioxide,  252.  The  estimation  of  lead  by  precipitation  as  oxalate 
and  titration  with  potassium  permanganate,  254. 


CHAPTER  VIII. 

NITROGEN;  PHOSPHORUS;  ARSENIC;  ANTIMONY;  BISMUTH; 

VANADIUM. 

Nitrogen.  —  The  determination  of  nitrogen  liberated  by  action  of  sodium 
hypobromite  upon  ammonia  compounds  and  derivatives,  256.  The  estima- 
tion of  nitrates  by  expulsion  of  nitrogen  pentoxide  on  ignition,  256.  The 
estimation  of  nitrates  by  reduction  with  a  ferrous  salt  and  titration  of  the 
residual  unoxidized  salt,  258.  The  estimation  of  nitrates  by  reduction  with 
ferrous  chloride  and  measurement  of  the  nitrogen  dioxide  evolved,  260.  The 
iodometric  determination  of  nitrates,  263;  action  of  manganous  chloride  in 
hydrochloric  acid,  263;  distillation  with  phosphoric  acid  and  potassium  iodide 
and  determination  of  iodine  in  the  distillate,  266;  decomposition  by  antimony 
trichloride,  determination  of  oxidation  of  residue  and  iodine  in  the  distillate, 
268.  The  iodometric  determination  of  nitrites,  269.  The  estimation  of  ni- 
trites, and  of  nitrites  and  nitrates  in  one  operation,  271;  determination 
of  nitrites,  271 ;  determination  of  nitrites  and  nitrates,  272.  The  estimation  of 
nitrates  and  chlorates  in  one  operation,  273.  The  qualitative  separation  and 


CONTENTS  i3t 

detection  of  ferrocyanides,  ferricyanides  and  sulphocyanates,  275;  the  ferro- 
cyanogen  ion,  275;  the  ferricyanogen  ion,  275;  the  ferrocyanogen  ion,  the 
ferricyanogen  ion  and  the  sulphocyanogen  ion  in  mixtures,  276.  The  gravi- 
metric determination  of  sulphocyanates,  276.  The  volumetric  estimation 
of  sulphocyanates  by  potassium  permanganate,  279.  Phosphorus.  —  The 
determination  of  phosphoric  acid  by  precipitation  as  ammonium  magnesium 
phosphate  and  weighing  as  magnesium  pyrophosphate,  282.  The  iodometric 
determination  of  phosphorus  in  iron,  283.  The  estimation  of  phosphoric  acid 
and  phosphorus  precipitated  as  ammonium  phospho-molybdate,  285.  The 
determination  of  phosphoric  acid  by  precipitation  as  uranyl  phosphate  and 
estimation  of  the  uranium  volumetrically,  286.  Arsenic,  Antimony,  and  Tin. 
—  The  determination  of  arsenic  by  precipitation  as  ammonium  magnesium 
arsenate  and  weighing  as  magnesium  pyroarsenate,  288.  The  iodometric  esti- 
mation of  arsenic  acid,  291;  reduction  by  hydriodic  acid  and  oxidation  by 
iodine  in  alkaline  solution,  291;  reduction  by  hydriodic  acid  and  titration  of 
iodine  liberated,  295.  The  detection  and  approximative  estimation  of  minute 
quantities  of  arsenic  in  copper,  301.  The  separation  of  arsenic  from  copper  by 
precipitation  as  ammonium  magnesium  arsenate,  305.  The  iodometric  deter- 
mination of  antimonic  acid,  and  of  antimonic  acid  and  arsenic  acid,  308.  The 
separation  of  antimony  from  arsenic  by  the  simultaneous  action  of  hydrochloric 
acid  and  hydriodic  acid,  and  the  estimation  of  antimony  in  the  residue,  311. 
The  detection  of  arsenic,  and  of  antimony  with  tin,  in  mixtures  containing 
compounds  of  these  elements,  312;  action  of  hydrochloric  acid  and  potassium 
iodide,  313;  action  of  hydrochloric  acid  and  potassium  bromide,  316.  The 
iodometric  determination  of  arsenic  and  antimony,  and  associated  copper,  318. 
The  estimation  of  arsenic,  antimony,  and  tin  in  the  lower  condition  of  oxida- 
tion by  the  action  of  potassium  ferricyanide  in  alkaline  solution  and  potassium 
permanganate  in  acid  solution,  322;  determination  of  antimony  in  anti- 
monious  condition,  323;  determination  of  tin  in  stannous  condition,  323; 
determination  of  arsenic  in  arsenious  condition,  324.  The  estimation  of  arsenic 
acid  and  antimonic  acid  associated  with  vanadic  acid,  325.  Vanadium.  — 
The  gravimetric  estimation  of  vanadic  acid  based  on  liberation  of  iodine  and 
absorption  of  that  element  by  silver,  325.  The  precipitation  of  ammonium 
vanadate  by  ammonium  chloride,  326.  The  estimation  of  vanadium  as  silver 
vanadate,  328.  The  estimation  of  vanadic  acid  by  the  action  of  the  halogen 
acids,  330;  the  action  of  hydrochloric  acid,  330;  the  action  of  hydrobromic 
acid,  335;  the  action  of  hydriodic  acid,  337.  The  determination  of  vanadic 
acid  by  reduction  in  acid  solution  and  reoxidation  by  iodine  in  alkaline  solu- 
tion, 341;  reduction  by  organic  acids,  341;  reduction  by  hydriodic  acid,  343; 
reduction  by  hydrobromic  acid,  345.  The  use  of  the  Jones  reductor  in  the 
estimation  of  vanadic  acid,  346;  regulation  of  reduction  by  the  use  of  silver 
sulphate,  348;  registration  of  reduction  by  use  of  ferric  sulphate,  349.  The 
estimation  of  vanadic  and  arsenic  acids  and  of  vanadic  and  antimonic  acids 
in  presence  of  one  another,  350.  The  estimation  of  vanadic  acid  associated 
with  chromium,  with  molybdenum,  and  with  iron,  352.  The  estimation  of 
vanadium  in  the  tetroxide  condition  by  the  action  of  potassium  ferricyanide 
in  alkaline  solution  and  potassium  permanganate  in  acid  solution,  352. 


CONTENTS 

CHAPTER  IX. 
OXYGEN;   SULPHUR;  SELENIUM;  TELLURIUM. 

Oxygen.  —  The  iodometric  determination  of  oxygen  in  air  and  in  aqueous 
solution,  355;  determination  of  oxygen  in  air,  355;  determination  of  dissolved 
oxygen,  360.  The  estimation  of  oxidizers  by  the  gravimetric  determination  of 
iodine  set  free  in  reaction,  361;  potassium  permanganate,  hydrogen  dioxide, 
potassium  dichromate,  ferric  chloride,  362.  Sulphur.  —  The  detection  of  sul- 
phides, sulphates,  sulphites,  and  thiosulphates  in  presence  of  one  another,  363. 
The  iodometric  determination  of  thiosulphates,  364.  The  iodometric  deter- 
mination of  sulphites  in  alkaline  solution,  366.  The  determination  of  dithionic 
acid  and  dithionates,  369.  The  determination  of  persulphates,  370;  arsenate- 
iodide  method,  370;  method  of  Le  Blanc  and  Eckardt,  371 ;  method  of  Grutz- 
ner,  372;  method  of  Mondolfo,  374;  method  of  Namias,  374.  Selenium. — 
The  gravimetric  estimation  of  selenious  acid  by  liberation  of  iodine  and  ab- 
sorption of  that  element  by  silver,  375.  The  gravimetric  determination  of 
selenious  acid  by  precipitation  of  selenium,  376.  The  iodometric  determina- 
tion of  selenious  acid  by  methods  based  upon  the  action  of  potassium  iodide 
in  presence  of  acid,  377;  the  contact  method,  377;  the  distillation  method,  379; 
the  differential  method,  treatment  of  the  residue,  380.  The  determination  of 
selenious  acid  by  potassium  permanganate,  382.  The  determination  of  sele- 
nious acid  by  the  direct  action  of  sodium  thiosulphate,  according  to  the  method 
of  Norris  and  Fay,  383.  The  iodometric  determination  of  selenic  acid  by  the 
action  of  the  halogen  acids,  385;  reduction  by  hydrochloric  acid,  with  distilla- 
tion, 385;  reduction  of  hydrobromic  acid,  with  distillation,  386;  reduction  by 
hydriodic  acid,  with  distillation,  388;  reduction  by  hydriodic  acid,  differential 
method,  388.  The  separation  of  selenium  from  tellurium  by  procedure  based 
upon  the  difference  in  volatility  of  the  bromides,  390.  Tellurium.  —  The 
gravimetric  estimation  of  tellurous  acid  by  the  liberation  of  iodine  and  absorp- 
tion of  that  element  by  silver,  394.  The  determination  of  tellurous  acid  by 
oxidation  with  potassium  permanganate,  394;  oxidation  in  presence  of  a  chlo- 
ride, 396;  oxidation  in  presence  of  a  bromide,  397.  The  determination  of 
tellurous  acid  by  the  precipitation  of  tellurous  iodide,  398.  The  iodometric 
estimation  of  tellurous  acid,  399.  The  iodometric  determination  of  telluric 
acid,  401.  The  precipitation  of  tellurium  dioxide  and  the  separation  of  tel- 
lurium from  selenium,  402. 

CHAPTER  X. 
CHROMIUM;  MOLYBDENUM;  URANIUM. 

Chromium.  —  The  estimation  of  chromium  as  silver  chromate,  406.  The 
iodometric  determination  of  chromic  acid,  407.  The  iodometric  estimation 
of  chromic  acid  and  vanadic  acid,  409.  The  estimation  of  chromic  acid  and 
vanadic  acid  by  reductions  and  oxidations,  411.  The  volumetric  estimation 
of  chromium  in  the  chromic  condition,  413.  Molybdenum.  —  The  gravimetric 
estimation  of  molybdic  acid  by  liberation  of  iodine  and  absorption  of  that 
element  by  silver,  414.  The  iodometric  estimation  of  molybdic  acid,  415;  the 


CONTENTS  XI 

digestion  method,  415;  distillation  process,  416;  reoxidation  of  the  residue  by 
iodine,  420;  reoxidation  of  the  residue  by  permanganate,  421.  The  estima- 
tion of  molybdic  acid  reduced  in  the  Jones  reductor,  424.  The  determination 
of  molybdic  acid  and  vanadic  acid  by  reductions  and  oxidations,  427.  Ura- 
nium. —  The  determination  of  uranium  by  the  aid  of  the  Jones  reductor,  430. 

CHAPTER  XI. 
FLUORINE;  CHLORINE;  BROMINE;  IODINE. 

Fluorine.  —  The  detection  of  fluorine,  432.  The  acidimetric  estimation  of 
fluosilicic  acid,  432.  The  iodometric  estimation  of  fluosilicic  acid,  435.  The 
estimation  of  fluorine  evolved  as  silicon  fluoride,  436;  elimination  of  silicon 
fluoride  at  high  temperatures,  436;  iodometric  determination  of  fluorine  in 
fluorides,  439.  Chlorine,  Bromine,  Iodine.  —  The  detection  of  iodine,  bromine, 
and  chlorine  in  presence  of  one  another,  440.  The  determination  of  free 
chlorine  and  free  bromine  by  liberation  of  iodine  and  absorption  of  that  element 
by  silver,  443.  The  gravimetric  determination  of  iodine  by  absorption  by 
metallic  silver,  444;  free  iodine,  444;  iodine  in  iodides,  446.  The  determina- 
tion of  halogens  in  benzol  derivatives  by  the  use  of  metallic  potassium,  447. 
The  direct  determination  of  chlorine  in  mixtures  of  alkali  chlorides  and  iodides, 
449;  use  of  ferric  sulphate,  449;  the  nitrite  method,  451.  The  direct  deter- 
mination of  bromine  (and  chlorine)  in  mixtures  of  alkali  bromides  (and  chlo- 
rides) with  iodides,  452.  The  application  of  iodic  acid  to  the  analysis  of  iodides, 
454.  The  iodometric  determination  of  iodine  in  haloid  salts,  457.  The  deter- 
mination of  the  halogens  by  the  electrolytic  reduction  of  silver  in  mixed  silver 
salts,  459;  silver  chloride  and  silver  bromide,  460;  silver  iodide  by  itself,  and 
in  mixture  with  silver  chloride  or  silver  bromide,  461.  The  estimation  of 
chlorates  by  reduction  with  ferrous  sulphate,  462.  The  iodometric  estimation 
of  chlorates,  463.  The  detection  of  alkali  perchlorates  associated  with  chlo- 
rides, chlorates,  and  nitrates,  465.  The  iodometric  determination  of  perchlo- 
rates, 467.  The  estimation  of  bromates  by  reduction  with  ferrous  sulphate, 
471.  The  iodometric  estimation  of  bromates,  471;  reduction  by  hydriodic 
acid,  471;  reduction  by  arsenious  acid,  474;  reduction  by  arsenate-iodide 
mixture,  475. 

CHAPTER  XII.  ' 
MANGANESE;   NICKEL;   COBALT;  IRON. 

Manganese.  —  The  determination  of  manganese  as  the  sulphate,  477.  The 
determination  of  manganese  as  oxide,  478.  The  determination  of  manganese 
separated  as  the  carbonate,  481.  The  determination  of  manganese  precipi- 
tated as  ammonium  manganese  phosphate  and  weighed  as  manganese  pyro- 
phosphate,  482.  The  electrolytic  determination  of  manganese,  485.  The 
determination  of  manganese  precipitated  by  the  chlorate  process,  487.  Nickel 
(Cobalt).  —  The  electrolytic  determination  of  nickel  with  the  rotating  cathode, 
489.  The  estimation  of  nickel  by  precipitation  as  the  oxalate  and  titration 
with  potassium  permanganate,  490.  The  detection  of  nickel  in  presence  of 


XU  CONTENTS 

cobalt,  491.  The  separation  of  nickel  and  cobalt  by  the  etherial  solution  of 
hydrochloric  acid,  492.  Iron.  — The  determination  of  iron  in  the  ferric  state 
by  reduction  with  sodium  thiosulphate  and  titration  of  the  excess  of  the  latter 
with  .  >dine,  492.  The  standardization  of  permanganate  in  iron  analysis,  495. 
The  benavior  of  ferric  chloride  in  the  Jones  reductor,  497.  The  effect  of 
nitric  acid  in  the  titration  of  a  ferrous  salt  by  potassium  permanganate,  498. 
The  permanganate  estimation  of  iron  in  presence  of  titanium,  499.  The  esti- 
mation of  iron  by  potassium  permanganate  after  reduction  with  titanous  sul- 
phate, 502.  Separations  of  iron  by  volatilization  in  gaseous  hydrogen  chloride, 
504;  iron  and  aluminium,  506;  iron  and  beryllium,  507;  iron  and  chromium, 
507;  iron  and  zirconium,  508.  The  estimation  of  iron  and  vanadium  in 
presence  of  each  other,  508.  The  estimation  of  ferric  iron,  vanadic  acid,  and 
chromic  acid  in  presence  of  one  another,  510. 


METHODS  IN  CHEMICAL  ANALYSIS 


CHAPTER   I. 
APPLIANCES  AND  GENERAL  PROCEDURE. 

MECHANICAL  PROCESSES. 

The  Determination  of  Products  Gaseous  at  Ordinary  Temperatures 

by  Loss  of  Weight. 

Various  forms  of  apparatus  have  been  designed  for  determin- 
ing, by  loss  of  weight,  reaction  products  which  are  gaseous  at 
ordinary  temperatures,  but  many  of  these  are  cumbersome  or 
require  skill  in  glass  blowing  for  their  construction.  A  form  of 
apparatus  described  by  Kreider*  is  light  and  easily  | 

made  from  three  test  tubes,  modified  and  fitted  as 
shown  in  the  figure.  The  test  tube,  A,  serves  as  the 
reaction  chamber.  B  is  perforated  with  a  hole  about 
I  cm.  in  diameter  and  fits  tightly  within  A;  and  C, 
so  selected  that  it  fits  loosely  within  B,  is  drawn  out 
to  a  small  capillary  tube. 

When  the  apparatus  is  to  be  used,  the  capillary 
of  C,  which  has  been  fitted  as  described,  is  pushed 
through  the  hole  of  B,  packed  loosely  with  cotton; 
B  is  filled  to  the  depth  of  from  6  cm.  to  8  cm.  (about 
two-thirds  of  its  contents)  with  granular  calcium  chlo- 
ride; and  B  and  C  are  adjusted  as  shown. 

To  the  test  tube,  C,  is  fitted  a  one-holed  stopper, 
through  which  passes  a  short  glass  tube  which  is  to  be 
closed  by  a  rubber  cap  and  plug.     Upon  removing  the 
plug,  and  applying  suction  to   the   short  tube,   the 
reagent  employed  to  liberate  the  volatile  product  to  be  deter- 
mined is  drawn  up  through  this  capillary  until  C  is  sufficiently 
filled.     Upon  replacing  the  plug  the  reagent  remains  within  C, 
held  by  atmospheric  pressure.      Gentle  pressure  upon  the  cap 

*  J.  Lehn  Kreider,  Am.  Jour.  Sci.,  [4],  xix,  188. 

i 


2  METHODS  TO  CHEMICAL  ANALYSIS 

expels  a  drop  of  liquid  from  the  capillary,  and  upon  the  release 
of  the  cap  a  little  air  is  drawn  in  to  allow  for  expansion  of  air 
in  the  large  tube  without  loss  of  liquid  during  subsequent 
handling. 

The  tubes  A  and  B  are  so  selected  that  very  little  of  the  product 
evolved  can  escape  between  them,  and,  in  case  they  fit  very 
loosely,  a  ring  of  paraffin  melted  into  the  mouth  of  A,  about  B, 
by  means  of  a  hot  wire,  seals  the  joint  securely.  A  very  con- 
venient way  to  attach  the  paraffin  is  to  melt  it  between  A  and 
another  tube,  which  fits  A,  as  does  B,  and  may  be  removed  by 
a  turning  motion,  leaving  the  ring  into  which  B  will  fit.  Very 
little  heating  is  then  required  to  make  a  tight  joint.  If  care  be 
used  in  taking  apart  A  and  B,  at  the  close  of  an  experiment,  such 
a  ring  of  paraffin  remains  in  place  and  may  be  used  many  times 
without  replacement,  being  remelted  by  a  touch  of  the  hot  wire 
before  every  new  experiment. 

In  making  a  determination,  the  substance  under  examination 
is  weighed  and  placed  in  the  bottom  of  A.  The  reagent  to  be 
employed,  10  cm.3  to  15  cm.3,  is  drawn  into  C,  and  held  there  in 
the  manner  described.  The  test  tube  A  is  slipped  over  B,  and 
the  joint  is  sealed  with  paraffin,  as  has  been  shown.  The  appa- 
ratus is  wiped,  placed  on  the  balance  and  weighed. 

Upon  removing  the  cap  from  the  small  tube  in  C,  the  reagent 
runs  from  C  into  A.  The  volatile  product,  forced  upward 
through  the  drying  column  of  calcium  chloride,  escapes  through 
the  annular  space  between  B  and  C.  When  action  ceases,  a 
current  of  dry  air  is  forced  through  C,  to  remove  all  the  volatile 
product,  the  cap  is  replaced,  and  the  apparatus  is  weighed.  The 
loss  of  weight  represents  the  volatile  product. 

Hydrogen  by  Loss. 


Metal  taken, 
grm. 

Hydrogen  found, 
grm. 

Error, 
grm. 

Magnesium  
Zinc 

O.  IOOO 
0.  IOOO 
O.  IOOO 
O.  IOOO 
O.  IOOO 
O.  200O 
O  .  2OOO 
O.  2OOO 
0  .  20OO 
O.  2OOO 

0.0087 
0.0085 
0.0084 
o  .  0084 
o  .  0083 
0.0061 
0.0062 
0.0062 
o  .  0060 
0.0061 

+0.0003 
-j-o.oooi 
o.oooo 

0.0000 
—  O.OOOI 

o  oooo 

+O.OOOI 

-j-o.oooi 

—  O.OOOI 

o.oooo 

APPLIANCES   AND   GENERAL   PROCEDURE  3 

Tests  of  this  apparatus  in  the  determination  of  carbon  dioxide 
in  carbonates,  and  of  nitrogen  in  urea  and  in  ammonium  salts, 
are  described  later. 

In  the  preceding  table  are  given, results  of  experiments  made  to 
determine  thus  the  weights  of  hydrogen  liberated  by  the  action 
of  magnesium  and  zinc  upon  dilute  hydrochloric  acid. 


The  Distillation  and  Condensation  of  Volatile  Products. 

The  rapid  evaporation  of  liquid  charged  with  soluble  or  in- 
soluble matter  is  apt  to  carry  mechanically  to  the  distillate 
some  material  which  should  remain  in  the  residue.  A  form 


Fig.  2. 

of  apparatus  elsewhere  described*  and  shown  in  the  accom- 
panying figure  (Fig.  2)  solves  the  problem  successfully.    The 
retort,  made  of  a  pipette,  bent  as  shown,  with  stoppered  funnel 
*  Gooch,  Am.  Chera.  Jour.,  ix,  28. 


4  METHODS   IN   CHEMICAL   ANALYSIS 

sealed  on  or  attached  by  a  rubber  joint,  is  fitted  to  an  upright 
condenser  which,  in  turn,  is  connected  by  a  stopper  to  a  thistle 
tube,  fitted  tightly  to  the  receiver  by  means  of  a  stopper  per- 
forated or  grooved  to  permit  the  passage  of  air.  For  work  to 
be  described  the  apparatus  has  been  modified  by  substituting 
for  the  perforated  or  grooved  stopper  a  tight  stopper  carrying 
a  bulbed  trap.* 

In  making  a  distillation,  the  liquid  is  introduced  by  the  funnel, 
the  glass  cock  is  closed,  the  water  started  through  the  condenser, 
and  the  retort,  not  more  than  half  filled  and  inclined  backward, 
is  carefully  heated.  For  the  heating  a  paraffin  bath  is  in  many 
cases  most  convenient,  and  it  is  advantageous  to  lower  the  retort 
into  the  paraffin,  already  heated  to  a  temperature  considerably 
above  the  boiling  point  of  the  liquid,  so  that  evaporation  may 
take  place  rapidly  and  often  without  actual  boiling.  The  diam- 
eter of  the  gooseneck  should  be  at  least  0.7  cm.  to  prevent  the 
formation  of  bubbles  within  it. 

The  use  of  this  apparatus  in  the  determination  of  boric  acid  is 
described  elsewhere. 

The  Distillation  and  Absorption  of  Volatile  Products. 

Apparatus  with  A  convenient  apparatus  for  the  distillation  and 
Ground  joints,  absorption  of  volatile  products  f  is  easily  constructed, 
with  glass  joints  throughout,  by  sealing  together  a  separating  fun- 
nel A,  a  Voit  flask  B,  a  Drexel 
wash  bottle  C,  and  a  bulbed  trap 
g,  as  shown  in  the  figure.  Upon 
the  side  of  the  distillation  flask 
B  is  pasted  or  etched  a  gradu- 
ated scale,  by  means  of  which 
the  volume  of  liquid  within  the 
flask  may  be  known  at  any  time. 
The  separating  funnel  is  con- 
.nected  with  a  Kipp  generator  set 
up  for  the  delivery  of  carbon  di- 
oxide, hydrogen,  hydrogen  chlo- 
FlS-  3-  ride  or  other  suitable  gas.  The 

flask  serves  as  the  retort,. the  wash  bottle  properly  charged  as 

*  See  trap  of  Fig.  7,  p.  6. 

t  F.  A.  Gooch  and  John  T.  Norton,  Jr.,  Am.  Jour.  Sci.,  [4],  vi,  168, 


APPLIANCES    AND    GENERAL   PROCEDURE 


5 


the  receiver,  and  the  products  of  distillation  are  swept  forward 
by  the  generator  gas,  which  may  serve  as  a  reagent  or  simply 
as  a  medium  for  aiding  the  transfer 
of  products  from  the  retort  to  the 
receiver. 

This  apparatus  has  served  a  useful 
purpose  in  processes  to  be  described 
for  the  determination  of  molybdenum, 
vanadium,  and  iodine  liberated  from 
the  iodide-iodate  mixture  by  acids  free 
or  evolved. 

A  similar  device  adapted  to  double 
distillation  is  shown  in  Fig.  4. 
will  be  given  later.* 


Fig.  4- 
Application  of  this  apparatus 


Fig-  5- 

Apparatus  with  Fig.  5  shows  a  convenient  device  by  Ed  gar,  f  put 
Sealed  joints,  together  without  ground  joints,  for  the  distillation 
and  absorption  of  volatile  products.  The  distillation  retort,  simi- 
lar in  design  to  that  of  Fig.  2,  consists  of  a  modified  pipette, 

*  F.  A.  Gooch  and  A.  W.  Peirce,  Am.  Jour.  Sci.,  [4],  i,  181. 
f  Graham  Edgar,  Am.  Jour.  Sci.,  [4],  xxvii,  174. 


6  METHODS   IN  CHEMICAL  ANALYSIS 

with  the  inlet  tube  bent  upward  and  sealed  to  a  separator/ 
funnel  while  the  outlet  tube,  expanded  to  a  small  bulb,  is  bent 
upward  and  then  downward  to  enter  the  absorption  flask. 

A  slow  current  of  hydrogen,  or  other  suitable  gas,  is  made  to 
enter  at  the  bottom  of  the  retort  to  stir  the  liquid  so  that  a  very 
small  volume  may  be  distilled  without  danger  of  "  bumping." 

The  Removal   of  Volatile  Products  without  Loss  of  Non-volatile 

Material  Reserved  for  Treatment. 

In  processes  which  involve  the  elimination  of  a  volatile  reagent 
or  product  of  reaction  from  a  boiling  solution,  it  is  often  essential 
to  prevent  losses  by  spattering  or  by  mechanical  transfer  of  non- 
volatile material  in  the  steam.  In  many  such 
processes  the  simple  device  shown  in  the  figure  is 
effective  in  preventing  appreciable  error  by  loss.* 
A  flask,  preferably  of  the  Erlenmeyer  shape,  with 
a  broad  bottom,  permits  boiling  of  the  liquid  in  a 
shallow  layer  favorable  to  the  checking  of  explosive 
ebullition.  A  two-bulbed  trap,  made  by  cutting 
short  an  ordinary  calcium  chloride  drying  tube 
and  hung  with  the  large  opening  downward,  ob- 
structs the  steam  while  permitting  sufficient  relief 
of  pressure  and  thus  serves  to  catch  and  return  to  the  liquid 
particles  of  the  non-volatile  matter  thrown  upward. 

The  Prevention   of   Mechanical    Loss   of   Solution   in   Reactions 

Evolving  Gaseous  Products. 

The  danger  of  mechanical  loss  in  reactions  ac- 
companied by  effervescence  (as  in  the  neutraliza- 
tion of  carbonates  by  strong  acids)  or  by  formation 
of  spontaneously  volatile  product  (as  in  the  libera- 
tion of  iodine  to  be  subsequently  titrated)  may  be 
minimized  by  making  use  of  a  trapped  reaction 
chamber.  For  this  purpose  the  apparatus  shown 
in  the  figure  is  serviceable.!  It  consists  of  a  Drexel 
washing  bottle  with  a  separatory  funnel  sealed  to 
the  inlet  tube,  and  a  Will  and  Varrentrapp  absorp-  Flg'  7> 
tion  apparatus  joined  to  the  outlet  tube.  The  reaction  is  brought 

*  F.  A.  Gooch  and  P.  E.  Browning,  Am.  Jour.  Sci.,  [3],  xxxix,  197. 
f  F.  A.  Gooch  and  C.  F.  Walker,  Am.  Jour.  Sci.,  [4],  iii,  293. 


u 


APPLIANCES   AND    GENERAL    PROCEDURE  7 

about  by  admitting  the  appropriate  reagents  through  the  funnel 
tube  to  the  solution  to  be  acted  upon  in  the  cylinder,  so  that  all 
volatile  products  must  escape  through  the  properly  charged 
absorption  bulbs. 

The  Transfer  of  Liquids  and  Gases  under  Pressure. 

A  simple  form  of  force  pump,  with  Bunsen  valves  of  special 
construction,  has  been  described  by  Kreider.* 

Valve.  —  In  forcing  a  liquid  or  gas  indifferent  to  rubber  from 
one  vessel  to  another,  the  ordinary  Bunsen  valve  is  apt  to  collapse 
in  such  a  way  as  to  permit  a  back  flow.  Kreider  finds  that  a 
stout  glass  tube  of  desirable  size,  sealed  at  one  end  and 
drawn  out  with  an  opening  in  the  constriction,  as  indi- 
cated in  the  accompanying  figure,  and  a  piece  of  rubber 
tubing  containing  a  smooth  slit  placed  over  it,  makes  a 
valve  in  which  collapse  is  impossible.  A  valve  similar  in 
appearance  to  the  one  here  described  has  been  previously 
used ;  but  the  similarity  is  confined  to  the  appearance,  as 
will  be  evident  from  the  following  description:  The  con- 
striction should  not  be  greater  than  is  necessary  to  leave  a  small 
space  between  the  glass  and  the  rubber  when  the  latter  is  loosely 
drawn  over  it ;  but  it  should  be  long  enough  to  permit  a  slit  of  about 
a  centimeter's  length  in  the  rubber  to  close  tightly,  or  about  twice 
the  length  of  the  slit.  A  slit  one  centimeter  long  will  be  found  to 
open  under  very  slight  pressure,  and,  to  accomplish  its  purpose, 
it  is  only  required  to  close  sufficiently  for  the  external  pressure  to 
force  the  rubber  against  the  opening  in  the  tube.  This  opening 
should  be  carefully  rounded  and  a  little  higher  rather  than  any 
lower  than  the  surrounding  glass,  and  is  better  made  before  seal- 
ing the  end,  in  order  to  keep  the  tube  perfectly  straight.  The 
rubber  should  fit  tightly  about  the  larger  parts  of  the  glass  tube 
and  be  put  on  with  care  to  have  the  smoothly  cut  slit  straight, 
and  loose  enough  to  close  tightly.  If  the  slit  is  placed  about 
90°  from  the  opening  in  the  tube,  sufficient  space  will  remain  to 
permit  the  escape  of  the  gas  or  liquid,  but  the  moment  the  pres- 
sure outside  becomes  greater  than  that  within,  the  rubber  will 
be  pressed  tightly  over  this  opening  and  thus  a  return  made  im- 
possible. When  dry  the  valve  does  not  resist  high  pressure  per- 
fectly; but  when  wet,  or  better,  when  both  glass  and  rubber, 
*  D.  A.  Kreider,  Am.  Jour.  Sci.,  [3],  1,  132. 


8  METHODS  IN   CHEMICAL  ANALYSIS 

including  the  slit,  are  moistened  with  glycerin,  a  nearly  perfect 
vacuum  may  be  retained  for  several  days.  The  valve  thus 
lubricated  with  glycerin,  when  used  as  a  protection  in  an  am- 
monia wash  bottle,  will  prevent  absolutely  the  access  of  am- 
monia to  the  mouth,  and  if  made  according  to  the  directions  will 
act  with  very  little  pressure.  Placed  in  the  connection  between 
the  vacuum  flask  and  water  pump  ordinarily  used  in  filtration, 
it  has  been  found  a  valuable  check  on  the  valve  of  the  pump, 
and  when  the  latter  fails  this  device  prevents  the  back  flow  of 
water  into  the  filtrate.  In  processes  which  necessitate  the  use 
of  a  partial  vacuum,  this  valve  may  be  employed  to  hold  the 
vacuum  in  continual  readiness. 

Force  Pump.  —  By  adjusting  two  of  the  valves  just  described 
to  the  opposite  extremities  of  a  T-tube,  with  the  horizontal  limb 
enlarged  or  sealed  to  a  larger  tube  so  as  to  permit  the 
attachment  of  a  large  and  stout  piece  of  rubber  tubing 
closed  at  one  end,  as  shown  in  Fig.  9,  a  convenient  and 
powerful  little  force  pump  is  obtained.  A  stout  T-tube 
of  small-bore  is  cut  off  short  at  the  two  ends  at  right 
angles  to  one  another;  to  one  is  sealed  a  tube  just 
large  enough  to  permit  the  insertion  of  a  valve;  to  the 
other,  a  large  tapering  tube,  slightly  lipped  so  as  to 
hold  a  piece  of  rubber  tubing  firmly  and  allow  of  tying 
the  latter  if  necessary.  Of  the  third  end  of  the  tube, 
a  valve  like  that  shown  in  Fig.  8  is  made.  The  com- 
pressing rubber  should  not  be  of  greater  length  than 
Fig.  9.  the  hand  is  able  to  cover  completely,  and  may  be 
closed  with  a  glass  stopper  selected  to  fit  tightly. 
Providing  the  space  through  the  T-part  is  kept  at  a  minimum 
compared  with  that  of  the  compressing  rubber,  rapid  pumping 
will  be  found  possible  and  the  power  limited  only  by  the  strength 
of  the  user's  grip.  The  apparatus  may  be  quickly  constructed 
of  materials  always  at  hand.  Originally  it  was  made,  in  about 
fifteen  minutes,  of  a  T-tube  to  which  the  necessary  enlargements 
were  connected  by  rubber  tubing  and  the  unused  space  filled  by 
a  glass  rod.  The  valves  may  be  inserted  directly  into  the  ends 
of  the  compressing  rubber,  but  the  form  shown  in  the  figure  is 
more  serviceable.  By  attaching  the  lower  end  to  a  tapering 
tube  as  shown,  the  pump  is  easily  inserted  into  a  perforated 
stopper  of  any  size. 


APPLIANCES   AND    GENERAL  PROCEDURE  9 

The  pump  has  been  found  serviceable  in  various  applications. 
For  filling  burettes  it  is  better  than  a  siphon,  the  stoppers  of  the 
standard  solution  bottles  being  provided  with  two  holes,  through 
one  of  which  the  delivery  tube  passes,  while  to  the  other  the 
pump  is  applied  by  the  adapter  shown  in  the  figure.  It  may  be 
applied  to  a  Kipp  generator  in  which  higher  pressure  is  momen- 
tarily required.  In  various  other  ways  it  has  been  found  to  be 
a  useful  piece  of  apparatus. 

A  Convenient  Form  of  Rotary  Shaker. 

An  apparatus  designed  by  Perkins*  serves  admirably  for  put- 
ting a  liquid  into  rotary  motion  for  the  purpose  of  securing 
gentle  but  thorough  agitation.  The  container  is  an  Erlenmeyer 


Fig.  10. 

flask.  This  is  suspended  in  a  retort  clamp  held  loosely  in  another 
clamp,  which  in  turn  is  also  loosely  held  by  another  clamp  firmly 
attached  to  the  upright  rod,  the  whole  forming  a  system  of  loose 
joints  at  right  angles,  which  permits  oscillatory  movement  of  the 
flask.  Motion  is  given  to  the  flask  by  a  wire  crank  attached 
eccentrically  to  the  rotating  table  driven  by  the  motor.  The  use 
of  this  apparatus  in  the  absorption  of  iodine,  free  or  liberated  in 
reaction,  by  metallic  silver  will  be  described  later. 

*  Claude  C.  Perkins,  Am.  Jour.  Sci..  [4],  xxviii,  33. 


10  METHODS   IN   CHEMICAL   ANALYSIS 

The  Purification  of  Precipitates  by  Solution  and 
Reprecipitation. 

In  many  processes  of  analytical  chemistry,  the  preparation 
of  substances  in  pure  condition  is  brought  about  by  precipita- 
tion, solution  and  reprecipitation ;  and  sometimes  this  cycle  of 
operations  must  be  repeated.  When  a  precipitate,  gathered 
upon  a  filter,  is  easily  acted  upon  by  the  appropriate  solvent, 
the  process  of  dissolving  the  precipitate  from  the  filter  is  simple; 
but  when  the  precipitate  is  refractory  toward  solvents  or  difficult 
to  attack  on  account  of  its  physical  condition,  as  is  the  case  with 
many  gelatinous  precipitates,  the  proper  handling  of  the  precipi- 
tate involves  some  inconvenience  and  delay. 

In  meeting  such  difficulties,  it  is  advantageous  to  place  within 
the  ordinary  paper  filter,  before  filtering,  a  movable  lining  of 
platinum  gauze  upon  which  the  precipitate  rests 
for  the  most  part  and  with  which  it  may  be  re- 
moved.* The  simplest  form  of  this  device  is 
easily  made  by  cutting  platinum  gauze  to  the 
shape  shown  in  the  accompanying  figure.  In 
ordinary  use,  this  piece  of  gauze,  folded  to  make 
a  cone  of  a  little  less  than  60°,  and  held  by 
pincers  at  the  point  of  overlapping,  is  placed 
Fig.  ii.  within  this  filter  and  allowed  to  fit  itself  closely 

by  the  natural  spring  of  the  gauze  when  released. 
Upon  filters  so  prepared  a  precipitate  may  be  collected  and 
washed  as  usual;  and,  at  the  end  of  the  operation,  the  cone  with 
nearly  all  the  precipitate  may  be  transferred  (conveniently  by 
means  of  ivory-pointed  pincers)  to  a  dish  or  beaker  for  suitable 
treatment.  The  small  amounts  of  the  precipitate  which  have 
passed  through  the  gauze,  being  somewhat  protected  by  the 
gauze  against  the  compacting  action  of  filtration  and  washing, 
are  generally  removable  with  ease  from  the  filter  by  a  jet  of  the 
washing  liquid.  After  washing,  the  gauze  may  be  replaced 
within  the  same  filter  and  serve  for  a  second  collection  of  the 
precipitate,  to  be  subsequently  dissolved,  in  case  double  precipi- 
tation and  solution  are  desirable.  The  final  collection  of  the 
precipitate  is,  of  course,  made  upon  paper  without  the  gauze  lin- 
ing, when  precipitate  and  filter  are  to  be  ignited. 
*  Gooch,  Am.  Jour.  Sci.,  [4],  xx,  n. 


APPLIANCES   AND   GENERAL  PROCEDURE  II 

This  device  has  proved  to  be  very  serviceable  in  handling  such 
precipitates  as  ferric  hydroxide,  aluminium  hydroxide  and  the 
basic  acetates. 

Precipitates  collected  upon  asbestos  in  the  perforated  crucible 
are  frequently  removable  without  difficulty  by  allowing  a  suitable 
solvent  to  percolate  precipitate  and  felt;  but  in 
case  the  precipitate  is  pasty  solution  in  this  man- 
ner may  be  unpleasantly  slow.  In  such  cases,  it 
is  convenient  to  remove  the  greater  part  of  the 
precipitate,  collected  and  washed  in  the  usual 
manner,  upon  a  disk  of  platinum  foil,  perforated, 
fitted  with  a  wire  handle,  as  shown  in  the  figure, 
and  placed  upon  the  asbestos  felt  before  the 
transfer  of  the  precipitate  to  the  crucible.  To 
make  such  a  disk,  shown  in  Fig.  12,  is  the  work 
of  a  few  moments  only;  and  by  its  use  pasty  Fi 
precipitates,  such  as  cuprous  sulphocyanate  or 
the  sulphides  of  the  metals,  are  easily  handled  for  solution. 


ELECTROLYTIC  PROCESSES. 
The  Rotating  Cathode. 

The  rotating  cathode,  previously  utilized  in  the  arts  in  the 
manufacture  of  seamless  copper  tubing  (by  electroplating  with 
currents  of  low  electromotive  force  and  continuous  replenish- 
ment of  the  bath  by  the  use  of  a  soluble  copper  anode) ,  and  by 
von  Klobukow*  for  slow  stirring  of  the  electrolytic  bath  in  analy- 
sis, has  been  applied  in  rapid  motion  by  Gooch  and  Medwayf- 
to  analytical  processes  in  which  the  object  is  to  remove  metals 
completely  and  expeditiously  from  solution.  An  ordinary  20  cm.3; 
platinum  crucible  is  used  as  the  cathode,  and  this  is  rotated  at  a 
speed  of  from  600  to  800  revolutions  by  means  of  a  small,  inexpen- 
sive electric  motor  fastened  so  that  its  shaft  is  vertical.  Upon 
this  shaft  the  crucible  is  fixed  by  pressing  it  over  a  rubber  stopper 
bored  centrally  and  fitted  tightly  on  the  end  of  the  shaft.  To 
secure  electrical  connection  between  crucible  and  shaft,  a  narrow 
strip  of  sheet  platinum  is  soldered  to  the  shaft  and  then  bent 

*  Jour,  prakt.  Chem.  (N.  F.),  xxxiii,  473. 

f    F.  A.  Gooch  and  H.  E.  Medway,  Am.  Jour.  Sci.,  [4],  xv,  320 


12 


METHODS   IN   CHEMICAL   ANALYSIS 


upward  along  the  sides  of  the  stopper,  thus  putting  the  shaft  in 
contact  with  the  inside  of  the  crucible  when  the  last  is  pressed 
over  the  stopper.  The  shaft  is  made  in  two  parts  as  a  matter  of 

convenience  in  removing  the  cruci- 
ble, and  is  joined,  with  care  to  make 
a  good  contact  between  the  two 
pieces  of  shafting,  by  a  rubber  con- 
nector of  sufficient  thickness  to  pre- 
vent the  crucible  from  wabbling 
when  rotated. 

The  solution  to  be  electrolyzed  is 
placed  in  a  beaker  upon  a  small  ad- 
justable stand,  so  that  the  crucible 
may  be  dipped  into  the  liquid  to  any 
desired  depth.  A  platinum  plate  is 
employed  as  an  anode,  and  this  is 
connected  to  the  positive  pole  of  a 
series  of  storage  batteries,  while  the 
negative  pole  of  this  series  is  con- 
nected to  the  bearing  in  which  the 
shaft  rotates,  thus  allowing  the  cur- 
rent to  go  from  the  platinum  plate 
through  the  solution  to  the  crucible,  up  the  shaft  of  the  motor, 
and  back  to  the  batteries.  The  power  to  run  the  motor  may  be 
conveniently  taken  from  the  incandescent  light  circuit. 

The  stand  carrying  the  beaker  is  raised  until  the  liquid  covers 
about  two-thirds  of  the  crucible  adjusted  to  the  shaft,  thus  giving 
a  cathode  surface  of  about  30  cm.2.  The  anode  is  introduced  and 
the  motor  started.  The  wires  from  the  storage  batteries  are 
connected  and  the  current  allowed  to  pass  through  the  solution. 
The  duration  of  the  electrolysis  is  varied  according  to  the  strength 
of  current  used,  but  in  each  case,  after  the  deposit  is  nearly  com- 
plete, the  current  from  the  batteries  is  shut  off,  the  motor  stopped, 
the  beaker,  platinum  anode  and  crucible  carefully  washed  with  a 
fine  jet  of  water,  the  motor  again  started,  and  the  current  allowed 
to  pass  for  the  remaining  time. 

When  the  deposition  is  complete  the  crucible  is  removed  and 
washed,  first  with  water,  then  with  alcohol,  and  finally  is  dried 
by  passing  it  over  a  flame. 

In  subsequent  study  of  the  material  and  shape  of  the  rotating 


Fig.  13. 


APPLIANCES   AND   GENERAL   PROCEDURE  13 

cathode,  Medway*  has  shown  that  a  silver  crucible  may,  with 
some  economy  and  without  sacrifice  of  accuracy,  be  substituted 
for  the  platinum  crucible,  at  least  in  the  determination  of  copper; 
that  neither  nickel  nor  aluminium  is  a  suitable  metal  for  use  as 
the  cathode;  and  that  a  rotating  disk  of  platinum  is  inferior  to 
the  crucible  for  use  as  the  cathode. 

The  results  of  experimental  tests  of  the  rotating  cathode  in  the 
determination  of  various  metals  —  copper,  silver,  nickel,  cadmium, 
tin,  gold,  zinc  —  are  given  under  the  headings  of  these  metals. 

The  Filtering  Crucible  in  Electrolytic  Analysis. 

The  rapidity  with  which  a  metal  or  oxide  may  be  thrown  upon 
the  electrode  and  thereafter  handled  successfully  in  the  ordinary 
processes  of  electrolytic  analysis  depends  upon  keeping  to  con- 
ditions under  which  deposits  are  compact  and  adherent.  It  is 
for  the  purpose  of  getting  adherent  deposits  that  in  modern  rapid 
processes  use  is  made  of  rotating  electrodes,!  of  apparatus  so 
arranged  that  gases  evolved  or  introduced  shall  stir  the  liquid, { 
and  of  the  agitating  action  of  a  magnetic  field. § 

The  use  of  these  methods  is,  however,  limited  to  those  cases 
in  which  attainable  conditions  and  the  nature  of  the  processes 
are  such  that  the  deposits  may  be  handled  and  washed  without 
loss  of  material  from  the  electrode.  Plainly,  the  range  of  con- 
ditions and  processes  may  be  very  much  extended  by  the  adop- 
tion of  means  for  handling  easily  and  safely  electrolytic  deposits 
more  or  less  loose.  Gooch  and  Beyer  ||  have  made  use  of  devices 
for  this  purpose,  in  which  the  filtering  crucible  of  platinum  or  of 
porcelain  is  adapted  to  use  as  an  electrolytic  cell. 
Electrolysis  and  Fig.  14  shows  a  convenient  form  of  apparatus  for 
Subsequent  such  use  in  electrolytic  analysis.  The  crucible  (A), 
fitted  in  the  usual  manner  with  an  asbestos  felt  (a), 
serves  as  an  electrode  (e)}  the  surface  of  which  is  very  much 

*  Am.  Jour.  Sci.,  [4],  xviii,  180. 

t  V.  Klobukow,  Jour,  prakt.  Chem.  (N.  F.),  xxxiii,  473.  Gooch  and  Med- 
way, Am.  Jour.  Sci.,  [4],  xv,  320.  (See  page  u.)  Exner,  Jour.  Am.  Chem.  Soc., 
xxv,  896. 

t  Levoir,  Zeit.  anal.  Chem.,  xxviii,  63.  Richards,  Jour.  Am.  Chem.  Soc., 
xxvi,  530. 

§  Frary,  Zeit.  Elektrochem.,  xiii,  308.     Jour.  Am.  Chem.  Soc.,  xxix,  1592. 

||  F.  A.  Gooch  and  F.  B.  Beyer,  Am.  Jour.  Sci.,  [4],  xxv,  249. 


METHODS    IN   CHEMICAL   ANALYSIS 


A 

F====== 


L 


Fig.  14- 


increased  by  a  layer  of  pieces  of  platinum  foil  (b)  within  the  crucible 
and  in  contact  with  its  walls.  The  joint  between  cap  and  crucible 
is  made  water-tight  by  a  thin  rubber  band 
(F).  The  capacity  of  the  cell  is  made  con- 
veniently ample  by  attaching  to  the  cruci- 
ble, by  means  of  a  close-fitting,  thin  rubber 
band  (E),  a  glass  chamber  (C)  easily  made 
from  a  wide,  short  test  tube.  The  second 
electrode  (/)  is  introduced  from  above 
through  the  glass  funnel  (D),  which  serves 
to  prevent  spattering  of  the  liquid  during 
the  electrolysis,  and  hangs  within  the  glass 
chamber.  The  cell,  held  by  a  clamp,  may 
be  kept  cool  during  action  by  immersing  it 
in  water  contained  in  a  cooler,  as  indicated 
in  Fig.  15. 

Electrical  connection  is  made  with  the 
crucible  by  means  of  a  platinum  triangle 
(c),  bent  as  shown  and  held  tightly  against 
the  outer  wall  of  the  crucible  by  a  rubber  band  (d).  Fig.  15 
shows,  on  the  left,  the  apparatus  adjusted  for  work. 

In  using  the  apparatus,  the  crucible,  fitted  with  asbestos  and 
containing  clippings  of  platinum  foil,  is  capped,  ignited  and 
weighed.  The  glass  chamber  with  the  wide  rubber  band  folded 
back  against  itself  is  set  upon  the  crucible  and  the  band  is  snapped 
into  place.  The  other  adjustments  are  made  in  the  manner 
shown.  The  electrolyte  is  introduced  and  the  current  turned  on. 
After  the  expiration  of  time  enough  to  complete  the  electrolysis, 
the  cooler  is  lowered  and  arrangements  are  made  to  draw  off  the 
liquid  in  the  cell.  If  the  process  is  such  that  no  harm  can  follow 
the  stopping  of  the  current  before  removing  the  liquid,  the  upper 
electrode  and  funnel  are  washed  and  removed,  the  cap  and  band 
are  slipped  off,  and  the  apparatus  is  set  in  the  holder  of  the  filter- 
ing flask  as  for  an  ordinary  filtration.  The  liquid  is  drawn 
through  the  felt  to  the  flask,  the  chamber  washed  down,  and  re- 
moved from  the  crucible,  and  the  deposit  is  well  washed.  The 
crucible  and  contents  are  dried  and  weighed,  the  increase  over 
the  original  weight  being,  of  course,  the  weight  of  the  deposit. 

The  details  of  experiments  made  to  test  this  form  and  use  of 
the  apparatus  are  given  in  the  table.  Copper  sulphate  strongly 


APPLIANCES   AND   GENERAL   PROCEDURE  15 

acidulated  with  sulphuric  acid  was  the  electrolyte.  Deposition 
was  completed  in  the  times  given,  and  the  ferrocyanide  test  ap- 
plied to  the  whole  nitrate  showed  the  absence  of  copper  in  every 
case.  The  apparatus  and  deposit  were  washed  first  with  water 
and  finally  with  alcohol.  It  was  noticed  that,  though  the  filtrate 
contained  no  copper,  the  washings  did  sometimes  contain  a  bare 
trace.  When  the  filtrate  was  allowed  to  stand  after  treatment 
with  potassium  ferrocyanide  it  turned  blue  rapidly,  and  this 


Fig.  is- 

action,  which  indicated  probably  the  presence  of  hydrogen  diox- 
ide or  of  persulphuric  acid  produced  in  the  electrolysis  o£  the  sul- 
phuric acid,  suggests  that  the  liquid  should  be  drawn  from  the 
deposit  as  quickly  as  may  be  after  the  current  is  cut  off.  In  the 
first  two  experiments  no  special  care  was  taken  in  this  respect, 
and  in  these  experiments  the  results  are  a  trifle  higher  than  those 
of  the  other  experiments,  in  which  the  manipulation  was  quickly 
made. 

Obviously  this  process  of  electrolytic  analysis  is  fairly  rapid, 
easily  executed,  and  accurate;    but  the  desirability  of  quickly 


i6 


METHODS  IN   CHEMICAL  ANALYSIS 


removing  the  liquid  from  the  deposit  after  stopping  the  current 
is  evident. 

Electrolysis  with  Filtration  after  Interruption  of  the  Current. 


CuSO4.5HjO 
taken. 

grm. 

Volume 
of  liquid. 

cm.1 

H2S04 
cm.* 

Current. 

Time, 
min. 

Theory  for 
copper. 

grm. 

Copper 
found. 

grm. 

Error, 
grm. 

Amp. 

Volt. 

0.5038 

50 

5 

i: 

5 
7 

41 

0.1283 

0.  1290* 

+0.0007 

0.5010 

50 

5 

U 

5 
7 

40  I 

o.  1276 

0.1282* 

+0  .  0006 

0.5009 

50 

&   s 

\24 

5 

7 

4! 

o.  1276 

0.1279* 

+0.0003 

0.5005 

50 

^'5 

ft 

5 

7 

0.1275 

0.1277* 

+  O.OOO2 

0.5047 

50 

f    5 

i: 

5 
7 

40  \ 

o.  1285 

0.1286* 

+0.0001 

0.5039 

50 

''    5 

i: 

5 

7 

,11 

^0.128^ 

,,  0.1285* 

+  O.OOO2 

o  .  5030 

50 

5 

5 

4 

o.  1281 

O.I282t 

+O.OOOI 

4 

7 

4°  ) 

*  No  copper  in  filtrate  or  in  washings. 


t  Trace  of  copper  in  washings. 


Electrolysis  Hollowing  are  results  obtained  as  in  the  preceding 

with  Filtration,  process  excepting  the  single  point  that  the  liquid  was 
drawn  off  while  the  current  was  still  running.  In  these  experi- 
ments the  nitration  was  effected  by  removing  the  cooler,  taking 
off  the  cap  and  band  from  the  crucible,  and  quickly  swinging  into 
place  the  filtration  apparatus  shown  at  the  right  in  Fig.  15.  The 
liquid  was  then  drawn  through  the  crucible  and  replaced  by  wash 

Electrolysis  and  Filtration  without  Interruption  of  the  Current. 


CuS04.sH»O 

taken. 

grm. 

Volume 
of  liquid. 

cm.3 

H2S04 
(i  :  i). 

cm.* 

Current. 

Time, 
min. 

Theory  for 
copper. 

grm. 

Copper 

found. 

grm. 

Error, 
grm. 

Amp. 

Volt. 

o  .  5030  * 

50 

5 

i; 

5 
7 

5} 

o.  1281 

0.1278! 

—  0.0003 

0.5008 

50 

5 

i: 

5 
7 

4! 

0.1275 

0.1275} 

0.0000 

0.5024 

50 

5 

i: 

5 

7 

40  1 

0.1280 

0.1277! 

-0.0003 

0.5014 

50 

5 

i: 

5 
7 

21 

0.1277 

o.  1276* 

—  O.OOOI 

0.5018 

50 

S 

t: 

5 

7 

40  } 

0.1278 

o  1278* 

o.oooo 

No  copper  in  filtrate  or  in  washings.  t  Trace  of  copper  in  filtrate. 

J  Trace  of  copper  in  washings. 


APPLIANCES   AND    GENERAL    PROCEDURE  IJ 

water  until  the  current  ceased  to  flow  because  there  was  no 
electrolyte  to  carry  it.  The  apparatus  was  washed  with  water 
and  finally  with  alcohol,  and  the  crucible  and  contents  were  dried 
for  periods  of  ten  minutes  at  ioo°-no°,  to  constant  weight. 

I.    When  a  deposit  is  so  loosely  adherent  as  to  be 

Electrolysis  ,.  ,  ,  , 

with  continuous  moved  by  the  liquid,  it  may  be  compacted  upon  the 
Filtration.  filtering  felt  by  keeping  the  liquid  in  process  of  filtra- 
tion and  constant  motion  through  the  cell  to  the  receiver.  The 
adjustment  of  apparatus  for  this  purpose  is  shown  in  Fig.  1 6. 


Fig.  16. 

Here  the  electrolytic  cell  rests  in  the  crucible  holder  fitted  to  a 
separating  funnel  used  as  a  receiver  and  connected  into  the  vac- 
uum pump.  A  stopcock  in  the  tube  of  the  crucible  holder  is 
convenient  but  not  necessary. 

The  manner  of  using  the  apparatus  is  simple.  First,  the 
weighed  crucible,  fitted  in  the  usual  manner  with  an  asbestos 
felt  and  containing  the  platinum  clippings,  is  adjusted  to  the 
glass  chamber.  The  cell  is  pressed  into  the  platinum  triangle 
and  set  into  the  holder.  The  funnel  which  carries  the  wire 


i8 


METHODS  IN  CHEMICAL  ANALYSIS 


electrode  is  put  in  place.  The  cell  is  charged  with  the  electrolyte 
and  the  current  is  turned  on.  The  electrolysis  begins  and,  under 
regulated  action  of  the  vacuum  pump,  the  liquid  is  drawn  through 
to  the  receiver  at  a  convenient  rate.  Usually,  before  the  upper 
electrode  is  uncovered  the  stopcock  is  closed,  the  suction  pump 
disconnected,  and  the  liquid  drawn  off  from  the  receiver  and 
returned  to  the  electrolytic  cell.  The  pump  is  again  connected, 
the  stopcock  is  opened  and  nitration  begins  again. 

Should  the  deposit  be  noticeably  loose,  it  may  be  compacted 
by  allowing  the  cell  to  drain  completely  under  the  action  of  the 
suction  pump.  The  electrolyte  is  thus  kept  in  circulation,  and 
loose  particles  of  the  deposit  are  held  upon  the  filtering  layer. 
From  time  to  time,  the  process  of  emptying  the  receiver  and 
filling  the  cell  is  repeated.  When  the  electrolysis  is  complete, 
as  shown  by  proper  testing  of  the  filtrate,  the  liquid  is  drawn 
through  the  crucible  and  replaced  by  water  from  above  until 
the  current  no  longer  flows.  The  electrodes  are  disconnected, 
the  extension  chamber  easily  slipped  off,  and  the  washing  of  the 
crucible  and  its  contents  continued  sufficiently,  with  care,  should 
the  deposit  be  spongy,  to  give  time  enough  in  the  washing  to 
properly  soak  out  absorbed  material.  The  crucible  and  contents 
are  dried,  ignited,  and  weighed  as  usual.  This  method  of  manipu- 
lation was  also  put  to  the  test  in  the  electrolysis  of  copper  sul- 
phate. Experimental  details  are  given  in  the  table. 

Electrolysis  with  Continuous  Filtration. 


CuSO4.sH2O 
taken. 

grm. 

Volume 
of  liquid. 

cm.3 

H2S04 
(i  :  i). 

cm.s 

Current. 

Time, 
min. 

Theory  for 
copper. 

grm. 

Copper 
found, 

grm. 

Error, 
grm. 

Amp. 

Volt. 

0-5013 

50 

5 

l: 

5 
7 

41 

0.1277 

O.I28of 

+0.0003 

0.5003 

50 

5 

): 

5 

7 

41 

0.1274 

0.1276! 

+O.OOO2 

0.5015 

50 

5 

i: 

5 
7 

4f 

0.1277 

0.1279* 

+O.OOO2 

0.5001 

50 

5 

i: 

5 

7 

si 

20) 

0.1274 

0.1274* 

O.OOOO 

0.5041 

50 

5 

i: 

5 

7 

41 

o.  1284 

0.1285! 

+  O.OOOI 

*  No  copper  in  filtrate  or  washings. 


t  Trace  of  copper  in  filtrate. 


J  Trace  of  copper  in  washings. 


The  results  show  that  there  is  no  difficulty  in  getting  accur- 
ate results  while  maintaining  continuous  filtration  during  the 


APPLIANCES   AND    GENERAL  PROCEDURE 


Fig.  17. 


process,  and  that  the  time  needed  to  complete  the  action  is 
somewhat  shortened  when  the  liquid  is  kept  in  circulation  by 
filtering. 

II.  Another  form  of  apparatus  for  electrolysis  with  continuous 
nitration,  in  which  a  porcelain  filtering 
crucible  replaces  the  platinum  filter  cru- 
cible, is  shown  in  Fig.  17.  In  this  appara- 
tus, it  is  necessary  to  make  the  connection 
from  above  with  the  electrode  inside  the 
crucible,  and  this  is  accomplished  by  a 
linked  platinum  wire,  as  shown.  In  put- 
ting together  and  using  this  apparatus,  a 
finely  perforated  disk  of  platinum  foil  (c) 
is  laid  upon  the  more  coarsely  perforated 
bottom  of  the  porcelain  crucible  (A). 
Upon  this  disk  the  asbestos  felt  (a)  is  de- 
posited in  the  usual  manner.  Platinum 
clippings  (b)  form  a  layer  of  suitable  thick- 
ness above  the  asbestos,  and  upon  this 
layer,  and  in  contact  with  it,  is  placed  another  perforated  disk 
of  platinum  foil  to  which  is  attached  a  twisted  wire  (e)  so  linked 
that  it  may  be  folded  within  the  crucible.  This  apparatus  is 
ignited  and  weighed,  and  to  it  is  adjusted,  as  shown,  a  cham- 
ber to  hold  the  electrolyte.  The  other  electrode  (/),  inclosed 
within  a  funnel  (D)  made  from  a  thistle  tube,  is  introduced 
in  the  manner  indicated.  This  apparatus  is  adapted  only 
to  use  in  the  method  of  continuous  filtration,  and  it  is  used 
exactly  as  in  Process  I.  Experimental  details  are  given  in  the 
table. 

By  either  of  the  processes  described,  reasonably  rapid  and 
accurate  electrolytic  determinations  may  be  made  without  the 
use  of  rotating  motors  or  special  stirring  apparatus,  and  without 
large  and  expensive  apparatus  of  platinum.  The  use  of  the  fil- 
tering crucible  as  a  part  of  the  electrolytic  cell  makes  possible 
the  utilization  of  operations  and  conditions  in  which  the  deposit 
may  lack  the  degree  of  adhesiveness  necessary  in  ordinary  electro- 
lytic processes.  The  application  of  the  processes  to  the  more 
difficult  determinations  of  manganese  and  lead  as  the  dioxides 
formed  upon  the  anode  in  very  imperfectly  adherent  condition 
will  be  described  later. 


20  METHODS  IN  CHEMICAL  ANALYSIS 

Electrolysis  with  Continuous  Filtration:   the  Use  of  the  Porcelain  Crucible. 


CuSO4.sH20 
taken. 

grm. 

Volume 
of  liquid. 

cm.3 

H2S04 
(i  :i). 

cm.3 

Current. 

Time. 

min. 

Theory  for 
copper. 

grm. 

Copper 
found. 

grm. 

Error. 

grm. 

Amp. 

Volt. 

(  2 

6 

5) 

0.5025 

50 

5 

<3 

8 

15 

o.  1280 

O.I277f 

—  0.0003 

(4 

10 

10  ) 

I    (2 

6 

5) 

0.5009 

50 

5              3 

8 

15 

o.  1276 

0.1279* 

+o  .  0003 

1  U 

10 

15) 

0.5025 

50 

6 

i: 

6 
10 

4! 

o.  1280 

0.1278* 

—  O.OO02 

0.5011 

50 

5 

i: 

6 

10 

41 

o.  1276 

o.i273f 

—  0.0003 

0.5013 

50 

5 

k 

6 
10 

41 

0.1277 

o.  1276$ 

—  o.oooi 

*  No  copper  in  filtrate  or  in  washings.  f  Trace  of  copper  in  filtrate. 

I  Trace  of  copper  in  washings. 

The  Fixation  of  Chlorine  on  the  Silver  Anode. 
Hydrochloric  From  a  consideration  of  the  apparently  very  exact 

Anode andPiati- resu^ts  obtained  by  many  investigators*  in  the  de- 
num  Cathode,  termination  of  the  chlorine  in  chlorides  by  fixation 
of  that  element  upon  an  anode  plated  with  silver,  it  would  seem 
that  nothing  could  be  simpler  than  the  accurate  determination 
of  the  chlorine  in  hydrochloric  acid  by  procedure  advocated  for 
the  treatment  of  metallic  chlorides.  That  such  is  not  the  case, 
however,  has  been  shown  by  Gooch  and  Read.f  In  an  experi- 
mental study  of  the  electrolysis  of  hydrochloric  acid  with  use 
of  the  silver  anode  and  platinum  cathode,  it  is  shown  that  silver 
oxide  is  formed  at  the  anode  and  must  be  decomposed  by  heating 
to  a  high  temperature  (incipient  redness,  at  the  tip  of  the  Bunsen 
flame),  not  simply  dried  over  a  steam  radiator |;  that  silver  de- 
posited from  the  double  cyanide  solution  upon  platinum  gauze  to 
make  the  silver  anode  always  includes  more  or  less  alkali  salt, 
which  is  lost  from  the  anode  surface  attacked  by  chlorine  during 
electrolysis;  that  to  avoid  contamination  of  the  silver  anode  by 
nonvolatile  material  it  should  be  plated  from  a  solution  of  silver 

*  Smith,  Jour.  Chem.  Soc.,  xxv,  890.  Myers,  ibid.,  xxvi,  1124.  Withrow, 
ibid.,  xxviii,  1350.  Hildebrand,  ibid.,  xxix,  447.  McCutcheon,  ibid.,  xxix, 
1445.  Lukens  and  Smith,  ibid.,  xxix,  1455.  Lukens  and  McCutcheon,  ibid., 
xxix,  1460. 

t  F.  A.  Gooch  and  H.  L.  Read,  Am.  Jour.  Sci.,  [4],  xxviii,  544. 

t  Smith's  Electro-analysis  (1907),  page  305. 


APPLIANCES  AND   GENERAL  PROCEDURE 


21 


oxalate  in  ammonium  hydroxide;  and  that  the  silver  anode  is 
attacked  and  dissolved  by  oxygen-chlorine  acids  produced  chiefly 
toward  the  end  of  the  electrolysis. 

The  Electrolysis  of  Hydrochloric  Acid  with  a  Silver  Anode  Plated  in  the  Oxalate 

Solution. 


Increase 

Apparent 

Apparent 

Chlorine 
taken  in 
HC1. 

<u 

6 
H 

Current. 

of  anode 
dried  in 
air  bath 
at 

Increase 
of  anode 
ignited. 

Loss  of 
dried 
anode  on 
ignition. 

error  in 
chlorine 
when 
anode 

error  in 
chlorine 
when 
anode  was 

I05°-no° 

was  dried. 

ignited. 

grm. 

min. 

Amp. 

Volt. 

grm. 

grm. 

grm. 

grm. 

grm. 

la 

0-0533 

25 

0.5  -O.OO 

2-4 

0.0513 

0.0502 

O.OOII 

—  O.OO2O 

—0.0031 

2b 

0.0533 

30 

0.45-0.01 

2-4 

0.0529 

0.0520 

o  .  0009 

—  0.0004 

—  0.0013 

W 

0.0533 

2=; 

0.5  -o.oi 

3-4 

0.0513 

0.0501 

O.OOI2 

—  O.OO2O 

—0.0032 

Ad 

0-0533 

30 

0.5  -o.oi 

3-4 

0-0523 

0.0517 

o  .  0006 

—  O.OOIO 

—  o  0016 

5« 

0-0533 

25 

0.5  -o.oi 

3-8-4 

0.0517 

o  .  0499 

O.OOlS 

—  0.0016 

—  0.0034 

6/ 

0-0533 

25 

0.4  -0.03 

3-8-4 

0.0502 

0.0493 

o  .  0009 

—0.0031 

—  O.OO4O 

7£ 

0-0533 

25 

0.4   -O.O2 

3-8-4 

o  .  0506 

0.0493 

0.0013 

—  0.0027 

—  O.OO4O 

a.  The  solution  became  suddenly  opalescent  and  soon  thereafter  the  current  practically  ceased, 
the  liquid  being  neutral  to  litmus  paper.     Silver  chloride  (0.0017  grm.)  was  recovered  from  the 
liquid,  and  silver  was  found  upon  the  cathode. 

b.  The  electrolysis  was  interrupted  at  the  first  appearance  of  opalescence,  the  liquid  being  neu- 
tral.    No  silver  was  found  in  solution  and  none  upon  the  cathode.    Iodine,  indicated  by  starch,  was 
liberated  when  potassium  iodide  was  added  to  a  portion  of  the  solution. 

c.  At  the  end  of  the  electrolysis  the  liquid  was  slightly  opalescent  and  neutral  to  litmus.    Upon 
the  addition  of  potassium  iodide  to  a  portion  of  it  a  trace  of  iodine  was  set  free.     In  another  portion, 
silver  nitrate  was  without  immediate  effect.     A  trace  of  silver  was  found  upon  the  cathode. 

d.  At  the  end  of  the  electrolysis  the  liquid  was  slightly  opalescent.     It  was  neutral  to  litmus, 
but  slowly  bleached  the  color.     Upon  standing  it  developed  acidity.     From  potassium  iodide  it 
liberated  iodine  equivalent  to  o.ooio  grm.  of  chlorine,  as  was  determined  by  sodium  thiosulphate. 
A  trace  of  silver  was  found  upon  the  cathode. 

e.  At  the  end  of  the  electrolysis  the  liquid  was  slightly  opalescent.     It  was  neutral  to  litmus 
but  developed  acidity  in  the  course  of  a  half-hour.    From  potassium  iodide  a  portion  set  free  iodine. 
In  the  remainder  silver  nitrate  gave,  after  standing  two  days,  an  amount  of  silver  chloride  equiva- 
lent to  0.0016  grm.  of  chlorine.     Silver  was  found  upon  the  cathode. 

/.  At  the  end  of  the  electrolysis,  the  liquid  was  slightly  opalescent.  It  was  neutral  to  litmus, 
but  developed  acidity  on  standing  a  half-hour.  From  potassium  iodide  a  small  portion  set  free 
iodine.  From  the  remainder  silver  nitrate  precipitated  silver  chloride,  which,  when  filtered  off 
after  five  days,  was  found  to  be  equivalent  to  0.0035  grm.  of  chlorine,  A  trace  of  silver  was  found  on 
the  cathode. 

g.  At  the  end  of  the  electrolysis,  the  liquid  was  slightly  opalescent.  It  was  neutral  to  litmus, 
but  developed  acidity  in  a  half-hour,  A  small  portion  of  it  set  free  iodine  from  potassium  iodide, 
Silver  nitrate  produced  in  the  remainder,  after  four  days,  a  precipitate  of  silver  chloride  equivalent 
to  0.0036  grm.  of  chlorine.  No  silver  was  found  upon  the  cathode, 

In  experiments  made  with  the  rotating  anode  of  gauze  plated 
with  silver  from  a  solution  of  silver  oxalate  in  ammonium  hydrox- 
ide, it  was  noted  that  near  the  end  of  electrolysis,  when  neutrality 
to  litmus  indicated  the  exhaustion  of  hydrochloric  acid,  the  solu- 
tion suddenly  became  opalescent  and  soon  afterward  the  current 


22  METHODS   IN   CHEMICAL  ANALYSIS 

practically  ceased  to  flow.  Upon  standing,  the  liquid,  which  at 
the  end  of  the  electrolysis  had  slowly  bleached  blue  litmus  paper 
without  reddening  it,  developed  distinct  acidity,  and  when  tested 
in  separate  portions  gave  further  opalescence  with  silver  nitrate 
and  set  free  iodine  from  potassium  iodide.  All  these  phenomena 
point  to  the  formation  of  hypochlorous  acid  in  the  process  of 
electrolysis  and  its  attack  upon  the  anode  to  form  silver  hypo- 
chlorite  and  derived  silver  salts.  It  appears  further  that  soluble 
silver  hypochlorite,  apparently  formed  chiefly  when  the  hydro- 
chloric acid  approaches  the  point  of  exhaustion,  is  thrown  into  so- 
lution, to  be  partially  decomposed  with  production  of  opalescent 
silver  chloride.  The  details  of  experiments  made  with  the  silver 
anode  plated  in  the  oxalate  solution  are  given  in  the  accompany- 
ing table.  In  the  first  experiment  the  electrolysis  was  continued 
until  the  current  practically  ceased  to  pass.  In  the  other  experi- 
ments the  operation  was  ended  when  the  diffusion  of  the  opal- 
escent silver  chloride  indicated  that  the  silver  anode  was  being 
attacked,  dissolved  and  partially  reprecipitated  in  the  liquid. 
In  all  the  liquid  was  neutral  at  the  end  of  the  electrolysis. 

The  results  of  these  experiments  show  the  fixing  of  oxygen  as 
well  as  chlorine  upon  the  anode,  the  removal  of  silver  from  anode 
to  cathode,  and  the  formation  of  hypochlorous  acid. 

It  is  plain,  therefore,  that  the  electrolytic  determination  of 
the  chlorine  of  hydrochloric  acid  with  the  use  of  the  silver  anode 
and  platinum  cathode  is  by  no  means  an  exact  process. 
Sodium  chloride,  The  electrolysis  of  sodium  chloride  with  the  use 
an'dM^cury  of  a  silver  or  silver-plated  anode  and  the  mercury 
cathode.  cathode  has  been  subjected  by  Peters*  to  a  very 

careful  and  critical  examination.  The  cell  used  was  similar  to 
that  described  and  figured  by  Hildebrand,f  consisting  of  a  bot- 
tomless beaker  6.3  cm.  in  diameter  and  6.3  cm.  tall  set  into  a 
crystallizing  dish  11.3  cm.  in  diameter  and  5.7  cm.  high.  Mid- 
way between  the  cells,  above  the  mercury,  was  a  coil  of  6  turns  of 
nickel  wire  I  mm.  in  diameter.  At  three  places  on  the  nickel 
wire  a  single  wire  ran  down  forming  legs  upon  the  feet  of  which 
rested  the  ends  of  a  Y.  This  Y,  made  of  glass  rod  3  mm.  in 
diameter,  formed  the  support  for  the  inner  cell  and  also  held  the 
nickel  wire  in  position  when  the  whole  apparatus  was  inverted  in 

*  Am.  Jour.  Sci.,  [4],  xxxii,  365. 
t  Jour.  Am.  Chem.  Soc.,  xxix,  451. 


APPLIANCES   AND    GENERAL   PROCEDURE 


emptying.  Three  rubber  stoppers  placed  radially  held  the  inner 
cell  in  position.  Mercury  sealed  the  two  compartments.  The 
contact  with  the  cathode  was  made  through  mercury  in  an 
S-shaped  tube,  hung  on  the  edge  of  the  outer  cell  with  a  platinum 
wire  sealed  in  one  end  and  dipping  under  the  cathode  mercury. 

The  anode  of  platinum  gauze 'or  of  pure  silver  gauze  was  rotated 
in  the  inner  cylinder. 

In  every  case  the  liquid  of  the  inner  cell,  charged  with  sodium 
chloride,  showed  alkalinity  to  indicators  whenever  tested.  In  a 
number  of  experiments  the  alkalinity  of  the  inner  cell  was  deter- 
mined by  titration  with  sulphuric  acid  using  methyl  orange. 
The  amounts  of  acids  used,  together  with  the  initial  and  final 
current  conditions,  are  given  in  the  table. 

Alkalinity  of  Inner  Cell. 


Sulphuric  acid, 
0.08996  N. 

cm.3 

Time  of 
electrolysis. 

min. 

Ampere,  initial 
and  final. 

Voltage,  initial 
and  final. 

Remarks. 

13.69 

140 

0.24-0.035 

4  cells 

8.66 

70 

0.  28-0.063 

4-3 

(  Changed  during 
\      experiment. 

2-39 

14 

1.4  -o.i 

7.2-8 

2.38 

18 

1.4  -O.I 

7.2-8 

2.15 

19 

1.3   -0.09 

7.2-8 

2.09 

18 

I  .  1  7-0  .  1 

7.4-8 

i-3'i 

42 

0.77-0.03 

4  cells 

I.  01 

38 

m^ 

6-8 

(  0.6  raised  to  i.o 
)    near  the  start. 

0.92 

*3 

o  .  63-0  .  03 

7.4-8 

0.62 

31 

1.3  -0.03 

7-8 

There  is  no  rapid  decrease  in  the  ammeter  reading  at  any  time 
to  indicate  the  complete  decomposition  of  sodium  chloride,  the 
accumulation  of  alkali  in  the  inner  cell  being  sufficient  to  carry  a 
considerable  amount  of  current. 

Silver  was  found  in  the  mercury  cathode  in  every  electrolysis 
with  anode  of  silver  or  silver-plated,  though  under  some  condi- 
tions the  amount  of  silver  transferred  to  the  cathode  was  small 
enough  to  be  negligible  in  analysis.  Upon  the  supposition  that 
the  transfer  of  silver  to  the  cathode  takes  place  in  greatest  degree 
toward  the  end  of  the  electrolysis  of  chloride,  the  experiment  of 
adding  the  sodium  chloride  in  small  portions  was  made  in  order 
to  keep  the  concentration  low  and  to  produce  repeatedly  in  one 
experimental  process  the  conditions  which  exist  toward  the  end 


METHODS  IN  CHEMICAL  ANALYSIS 


of  the  ordinary  electrolysis.  In  this  manner  the  maximum  trans- 
fer of  silver  for  a  given  amount  of  chloride  electrolyzed  should  be 
obtained.  The  details  of  experiments  made  in  this  manner,  and 
of  other  experiments  in  which  the  entire  amount  of  chloride  to 
be  electrolyzed  was  introduced  at  once,  are  shown  in  the  table. 
The  results  show  the  transfer  of  silver  from  the  plated  anode 
to  the  mercury  to  be  a  regular  feature  of  the  electrolysis,  under 
the  conditions  described,  and  that  the  amount  of  silver  so  trans- 
ferred is  considerably  increased  if  the  electrolysis  is  continued 
beyond  the  point  where  the  sodium  chloride  is  all  decomposed; 
but  that  if  the  electrolysis  is  interrupted  as  soon  as  the  salt  is  all 
decomposed,  very  little  silver  is  transferred  to  the  mercury. 

Transfer  of  Silver  to  the  Cathode. 


Total  NaCl 
sol.,  o.i  N, 
added. 

Number  of 
additions. 

Number  of 
times  anode 
was  heated 
during  process 

Total  time 
of  elec- 
trolysis. 

Initial  and  final  current 
conditions. 

Amount  of 
silver 
recovered 
by  distil- 
lation. 

cm.3 

to  free  from 
oxide. 

min. 

Amperes. 

Volts. 

grm. 

Silver-plated  anode. 


47.98 

48.12 

24 
24 

3 
3 

305 
232 

::£h- 

8 

\  0.0075 
\  0.0083 

Silver  anode. 

17* 

12 

4 

1077 

0.0162 

Silver-plated  anode. 

50.00 
50.00 

I 
I 

o 

0 

38 

31 

0.6  -0.3 
i.o  -0.3 

7.4-8 
7.4-8 

o  .  0006 
o  .  0005 

Silver-plated  anode. 

50.02 
50.00 
50.00 
50.00 

I 
I 
I 
I 

o 
o 

0 

o 

14 

18 

iQ 
18 

1.4  -o.i 
1.4  -o.i 
1.3  -0.09 
i  .  i  7-0  .  i 

7.2-81 
7.  2-8  I 
7.2-8  f 

7-4-8J 

0.0005* 

Silver-plated  anode. 


50 
50 
50 
50 

I 
I 
I 
I 

o 

0 

o 
o 

14 

16 

13 
ii 

1.2    -O  .  2 
1.2    -O  .  2 
1.5    -O.IQ 
I  .  2    -O.IQ 

7-4-8] 
7-4-8! 
7-4-8  | 

7-4-8J 

o.oooif 

*  Examination  of  the  liquid  of  the  inner  cell  showed  that  the  decomposition  of  sodium  chloride 
was  complete. 

t  Examination  of  the  liquid  of  the  inner  cell  showed  that  the  decomposition  of  sodium  chloride 
was  not  complete.  Amounts  varying  from  8  mgrm.  to  30  mgrm.  of  sodium  chloride  were  found  at 
the  end  of  the  operation. 


APPLIANCES  AND    GENERAL   PROCEDURE  25 

The  experimental  data  were  utilized  by  Peters  to  determine 
the  best  conditions  for  the  electrolysis  of  50  cm.3  of  n/io  sodium 
chloride  with  the  apparatus  at  hand.  The  procedure  prescribed 
for  the  quantitative  electrolysis  of  .2923  grm.  of  sodium  chloride 
in  50  cm.3  of  water  is  as  follows:  Introduce  about  2  kg.  of  mer- 
cury into  the  apparatus  or  sufficient  mercury  to  rise  6  mm.- 
8  mm.  above  the  bottom  of  the  inner  cell.  Cover  the  nickel  wire 
with  water  (70  cm.3-8o  cm.3)  and  add  to  it  I  cm.3  of  saturated 
salt  solution.  Introduce  the  50  cm.3  of  salt  solution  (.2923  grm. 
NaCl)  into  the  inner  cell  and  electrolyze  with  the  bottom  of  the 
anode  6  mm.-io  mm.  above  the  surface  of  the  mercury.  The 
current,  1.2  amp.-i.5  amp.  at  the  start,  should  be  stopped  at 
o.i  amp.  Take  out  the  anode,  hold  it  over  the  inner  cell  and 
wash  with  the  jet  of  the  wash  bottle.  Introduce  into  the  heat- 
ing apparatus  (the  electric  furnace  or  the  heating  crucible) ;  heat 
at  4OO°-45O°,  and  raise  the  temperature  during  5-10  minutes 
until  the  silver  chloride  is  fused,  avoiding  temperatures  over  550°. 
Cool  and  weigh.  Titrate  the  liquid  over  the  mercury  to  near 
the  end-point,  using  methyl  orange  with  n/io  hydrochloric  acid, 
then  transfer  to  a  separatory  funnel  and  shake  well  to  decompose 
the  remaining  sodium  amalgam.  Draw  off  the  mercury  and 
finish  the  titration,  using  sodium  hydroxide  to  determine  the 
end-point.  In  case  violent  shaking  introduces  suspended  mer- 
cury into  the  liquid,  making  the  end-point  difficult  to  determine, 
the  mercury  may  be  allowed  to  settle  out,  or  the  mixture  may  be 
filtered,  or  more  methyl  orange  may  be  added. 

The  results  of  experiments  made  in  accordance  with  the  pro- 
cedure outlined  are  given  below.  The  time  required  for  a  single 
determination  was  one  hour.  The  temperatures  given  are  in 
each  case  those  of  the  electric  oven  at  the  beginning  and  end  of 
heating. 

Under  the  most  favorable  conditions,  determined  empirically 
for  one  definite  amount  of  sodium  chloride,  results  show  an  aver- 
age of  fair  analytical  accuracy.  The  chlorine  tends  to  be  low, 
averaging  —0.0005  grm-  between  the  extremes  of  +0.0008  grm. 
and  —0.0030  grm.;  and  the  sodium  high,  averaging  +0.0005 
grm.  between  the  extremes  of  +0.0016  grm.  and  —0.0018  grm. 
What  the  proper  conditions  would  be  with  a  sodium  chloride 
solution  of  different  strength,  an  apparatus  of  slightly  different 
dimensions,  or  different  current  conditions,  the  author  feels  would 


26 


METHODS  IN  CHEMICAL  ANALYSIS 


have  to  be  determined  by  preliminary  experimentation.  The  am- 
meter cannot  be  relied  lupon,  without  previous  experimentation, 
to  determine  the  end  of  the  electrolysis,  as  the  current,  during 
the  operation,  falls  gradually  from  the  start  to  an  indefinite  end. 
There  is  no  rapid  decrease  in  the  ammeter  reading  at  any  point 
to  indicate  the  complete  decomposition  of  the  sodium  chloride, 
the  accumulation  of  the  .alkali  in  the  inner  cell  being  sufficient  to 
carry  a  considerable  current. 

Analysis  of  Sodium  Chloride. 


Chlorine 
found. 

grm. 

Sodium 
found. 

grm. 

Time  of 
electrolysis. 

min. 

Time  of 
heating 
anode. 

min. 

Temperature 
of  heating 
anode. 

Error, 
chlorine. 

grm. 

Error, 
sodium. 

grm. 

0.1755 

0.1152 

19 

IO 

I  45°l 

—  O.OOlS 

+O  .  OOO2 

0.1771 

O.II59 

21 

23 

J46o| 

—  0.0002 

+0.0009 

0.1770 

0.1158 

18 

14 

|5So| 

—0.0003 

+0.0008 

0.1781 
0.1775 

0.1166 
0.1157 

18 

20 

13 
19 

Uoo) 
I47o( 

isi 

+0.0008 

+O.OOO2 

+O.OOI6 
+0.0007 

0.1743 

0.1132 

2O 

72 

1  480! 

—  O.OO30 

—  0.0018 

0.1780 

0.1161 

19 

7 

{405 

1  440  $ 

+O.OOO7 

+O.OOII 

0.1767 

0-1155 

19 

6 

i  440} 

—  0.0006 

+0.0005 

Aver.  0.1768 

o.ii55 

19 

—  0.0005 

+0.0005 

The  work  shows  that  in  the  electrolysis  of  sodium  chloride, 
Under  the  conditions  described,  with  the  anode  of  silver,  or  silver- 
plated,  and  the  mercury  cathode,  silver  is  always  transferred 
from  the  anode  to  the  cathode  mercury;  although,  under  condi- 
tions determined  by  experiment  to  be  most  favorable  for  a  given 
amount  of  chloride,  the  amount  of  silver  transferred  may,  for 
analytical  purposes,  be  neglected. 

With  the  apparatus  described,  the  electrolysis  of  50  cm.3  of 
n/io  sodium  chloride,  0.2923  grm.  of  the  salt,  is  most  favorably 
accomplished  with  a  current  of  1.2-1.5  amp.  at  the  start,  which 
is  allowed  to  fall  off  to  O.I  amp.,  the  operation  requiring  18-20 
minutes. 


APPLIANCES   AND   GENERAL   PROCEDURE  27 

It  is  recommended  that  the  anode  covered  at  the  end  of  the 
electrolysis  with  silver  chloride  be  first  heated  below  the  fusing 
point  of  the  chloride  to  decompose  all  of  the  silver  oxide,  and 
then  for  five  or  ten  minutes  at  a  temperature  of  4OO°-5OO°  to 
fuse  the  chloride. 

Sodium  hydroxide  is  always  present  in  the  inner  cell  after  the 
beginning  of  the  electrolysis. 

The  best  method  of  treating  the  anode  covered  with  fused  silver 
chloride  to  prepare  it  to  be  used  in  a  subsequent  electrolysis  is 
by  heating  about  20  minutes  at  about  500°,  or  a  little  over,  in  a 
current  of  hydrogen. 

IODOMETRIC  PROCESSES. 

The  Standardization  of  Iodine  Solutions  by  the  Action  of  Metallic 

Silver. 

The  affinity  between  silver  and  iodine  has  been  made  by  Gooch 
and  Perkins*  the  basis  of  a  gravimetric  method  for  the  determi- 
nation of  iodine  in  general,  and,  incidentally,  for  the  gravimet- 
ric standardization  of  iodine  solutions  to  be  used  in  volumetric 
analysis.  In  studying  the  absorption  of  iodine  by  silver,  experi- 
ments were  made  with  silver  reduced  in  the  wet  way  by  the  action 
of  zinc  upon  silver  chloride,  nitrate  or  iodide;  in  the  dry  way,  by 
the  action  of  hydrogen  upon  silver  sulphide  or  oxide;  or,  electro- 
lytically,  from  a  solution  of  silver  nitrate,  with  anode  inclosed  in 
a  porous  cell  and  with  an  oscillating  cathode  of  platinum,  experi- 
ence having  shown  that,  while  the  bright  and  crystalline  deposit 
formed  upon  a  stationary  cathode  is  lacking  in  absorptive  power,, 
the  broken  and  dark  product  obtained  by  continually  oscillating 
and  scraping  the  cathode  during  the  deposition  of  the  metal  is 
sensitive  to  iodine.  Silver  reduced  from  a  silver  salt  by  zinc  or 
from  silver  sulphide  by  hydrogen  may  serve  the  purpose,  provided 
it  is  subjected  to  a  preliminary  treatment  with  potassium  iodide* 
and  silver  reduced  from  the  oxide  by  hydrogen  is  also  service- 
able; but  the  best  form  of  silver,  and  the  one  most  easily  pre- 
pared in  the  pure  state,  is  that  deposited  electrolytically  as  a 
black  mass  upon  a  small  oscillating  cathode  of  platinum  from 
a  solution  of  silver  nitrate,  the  platinum  anode  being  inclosed  in 
a  porous  cell.  As  long  as  the  mass  adheres  to  the  electrode  it 
*  F.  A.  Gooch  and  Claude  C.  Perkins,  Am.  Jour.  Sci.,  [4],  xxviii,  33. 


28  METHODS  IN  CHEMICAL  ANALYSIS 

remains  perfectly  intact,  but  as  soon  as  it  is  shaken  off  into  the 
solution  it  turns  to  a  dull-gray  color  and  settles  to  the  bottom 
in  a  fine  flourlike  powder.  It  should  not  be  allowed  to  remain 
upon  the  electrode  after  it  begins  to  change  color,  as  the  silver 
then  collects  in  a  crystalline  form  and  does  not  absorb  iodine 
readily. 

It  was  shown  that  when  finely  divided  silver  is  shaken  with  a 
solution  of  iodine  in  potassium  iodide,  the  iodine  is  absorbed  by 
the  metal,  but  that  more  iodine  is  absorbed  than  was  originally 
free  if  the  mixture  is  exposed  to  air  in  the  shaking.  Experiments 
in  which  the  silver  was  shaken  with  50  cm.3  of  a  solution  of  potas- 
sium iodide,  20  grm.  to  the  liter,  proved  fully  that,  either  in 
neutral  or  alkaline  solution,  the  action  of  air  must  be  prevented 
during  the  agitation  of  the  iodide  with  silver;  for  in  these  experi- 
ments it  was  found  that,  from  the  solution  of  potassium  iodide 
shaken  in  contact  with  air,  finely  divided  electrolytic  silver 
absorbed  o.ooio  grm.  of  iodine  in  fifteen  minutes,  silver  reduced 
by  zinc  from  silver  iodide  0.0012  grm.  in  fifteen  minutes,  silver 
reduced  from  the  sulphide  by  hydrogen  0.0032  grm.  in  one  hour, 
and  crystalline  electrolytic  silver  0.0051  grm.  in  one  hour  and 
forty-five  minutes. 

The  procedure  finally  shown  to  be  best  is  the  following:  Into 
a  25O-cm.3  Erlenmeyer  flask  are  put  the  iodine,  in  solution  in 
potassium  iodide,  and  a  weighed  amount  of  the  finely  divided 
silver.  The  glass  stopper  of  a  Drexel  bottle  is  fitted  into  the 
neck  of  the  flask  and  held  in  a  tight  joint  made  by  slipping  over 
neck  and  stopper  a  broad  thin  rubber  band.  The  air  in  the  flask 
is  replaced  by  hydrogen  and  the  tubes  of  the  stopper  are  capped. 
The  flask  thus  trapped  is  shaken  with  a  rotary  motion  by  hand, 
or  preferably  by  a  mechanical  shaker,  until  the  iodine  color  van- 
ishes. The  liquid,  usually  50  cm.3  in  volume,  is  diluted  to  about 
100  cm.3  and  the  residue  of  silver  and  silver  iodide  is  collected 
in  a  perforated  crucible  fitted  with  asbestos  felt,  washed,  dried 
at  130°  to  140°  and  weighed.  The  difference  between  the  weight 
of  silver  taken  and  that  of  the  residue  of  silver  and  silver  iodide 
is  the  measure  of  the  free  iodine.  Fig.  10  shows  the  adjustment 
of  the  apparatus. 

Figures  to  show  the  accuracy  of  the  procedure  are  given  in 
connection  with  processes  for  the  determination  of  iodine. 


APPLIANCES  AND   GENERAL   PROCEDURE  2£ 

Arsenic  Trioxide  as  an  Iodometric  Standard. 

Carefully  sublimed  and  anhydrous  arsenic  trioxide  serves  as 
the  most  exact  standard  in  iodometric  processes,  and  it  is  also  use- 
ful *  in  the  standardization  of  potassium  permanganate  for  quan- 
titative oxidations. 

To  make  a  standard  arsenite  solution  it  is  convenient  to  dissolve 
4.95  grm.  of  arsenic  trioxide  in  a  concentrated  solution  of  4  grm. 
of  potassium  hydroxide,  and  to  make  up  the  solution  to  I  liter  for 
the  nj  10  or  to  2  liters  for  the  n/2O  solution  by  adding  100  cm.3  or 
200  cm.3  of  a  saturated  solution  of  potassium  hydrogen  carbonate 
and  water  to  complete  the  volume.  The  arsenic  trioxide  dis- 
solves much  more  readily  in  the  alkali  hydroxide  than  in  the 
acid  carbonate  alone,  and  the  proportion  given  is  not  enough  to- 
entirely  form  dipotassium  hydrogen  arsenite,  but  more  than 
enough  to  form  potassium  dihydrogen  arsenite. f  Should  more 
alkali  hydroxide  be  used  in  effecting  solution,  it  should  be  neu- 
tralized by  suitable  amount  of  acid  (sulphuric  acid  or  hydro- 
chloric acid)  before  the  final  addition  of  the  acid  carbonate  ia 
large  excess;  but  it  is  important  to  obtain  the  neutralization 
without  the  addition  of  any  indicator  containing  alcohol,  and 
this  is  easily  accomplished  without  the  use  of  an  indicator  if  at- 
tention is  paid  to  the  amount  of  alkali  hydroxide  employed. 

With  this  standard  arsenite  solution  the  standardization  of 
iodine  in  potassium  iodide  is  effected  by  titration  with  or  without 
starch  as  an  indicator,  according  to  circumstances. 

The  Starch  Indicator  for  Free  Iodine. 

In  the  literature  of  iodometric  titration  frequent  mention  is 
made  of  the  production  of  a  red  color,  as  well  as  a  blue,  when 
starch  is  used  as  an  indicator,  and  numerous  formulae  have  been 
given  for  making  a  starch  solution  designed  to  give  a  blue  color 
with  iodine,  and  to  keep  without  spoiling.  An  investigation  by 
Hale  t  into  the  occasion  of  the  loss  of  iodine  met  with  in  titrimet- 
ric  processes  in  connection  with  the  formation  of  a  red  color, 
has  led  not  only  to  the  elimination  of  the  loss,  but  also  to  a 

*  See  page  41. 

t  Hale,  Am.  Jour.  Sci.,  [4],  xiii,  387. 

I  F.  E.  Hale,  Am.  Jour.  Sci.,  [4],  xiii,  379. 


METHODS  IN  CHEMICAL  ANALYSIS 


probable  explanation  of  the  cause  of  the  red  color,  of  the  loss  of 
iodine,  and  of  their  mutual  relation. 

Hale  observed  that,  in  titrating  a  decinormal  solution  of  arse- 
nite  with  a  decinormal  solution  of  iodine,  the  amount  of  iodine 
needed  was  greater  when  dependence  was  placed  upon  the  occur- 
rence of  the  starch  blue  following  intermediate  tints  of  red  to 
'determine  the  end-point,  than  when  the  titration  was  made  with- 
out the  indicator  to  the  appearance  of  the  yellow  tinge  due  to 
the  addition  of  a  single  drop  of  decinormal  iodine  in  excess. 

The  loss  of  iodine  in  the  interaction  of  decinormal  solutions  in 
presence  of  the  starch  indicator  used  in  a  certain  series  of  experi- 
ments is  shown  in  the  statement  below  to  be  very  considerable 
for  the  undiluted  solutions  and  much  less  considerable  when  the 
reaction  took  place  in  dilute  solution. 

The  Decinormal  Solutions  Undiluted. 


n/io  As2O3  sol. 

Reading  with  1.25  cm.3 
of  starch  solution  to 
a  deep  blue. 
Nearly  n/io  I. 

Iodine  reading 
calculated. 
Nearly  w/io  I. 

Error. 

cm.* 

cm.' 

cm.8 

cm.8 

5 

5-23 

5-14 

O.OQ 

IO 

10.37 

10.28 

O.Og 

IS 

IS-50 

I5-42 

0.08 

20 

20.70 

20.56 

0.14 

25 

25.85 

25.70 

0.15 

3° 

31.00 

30.84 

o.  16 

35 

36.12 

35.98 

0.14 

40 

41  .  2O 

41.08 

0.12 

45 

46.39 

46.26 

0.13 

50 

51-54 

5L40 

0.14 

Final  Volume  no  cm.3 


tt/io  As2O3  sol. 

Reading  with  1.25  cm.3 
of  starch  solution  to 
a  deep  blue. 
Nearly  w/io  I. 

Iodine  reading 
calculated. 
Nearly  M/IO  I. 

Error. 

cm.3 

cm.8 

cm.3 

cm.3 

I 

1-05 

1.03 

O.O2 

2 

2.10 

2.05 

0.05 

3 

3-15 

3.08 

O.O7 

5 

5-20 

5-14 

O.O6 

7 

7.25 

7.20 

0.05 

IO 

10.32 

10.28 

O.O4 

15 

15.47 

I5.42 

0.05 

20 

20.60 

20.56 

0.04 

35 

36.02 

35.98 

O.O4 

APPLIANCES  AND   GENERAL  PROCEDURE  31 

If  the  starch  solution  be  added  after  the  yellow  color  of  free 
iodine  is  visible  a  fine  blue  is  produced  by  a  preparation  which 
would  give  the  intermediate  colors  if  it. were  present  during  the 
entire  titration.  This  fact  shows  that  some  cause  for  the  pro- 
duction of  the  red  lies  in  the  titration.  J  It  excludes  any  explana- 
tion of  the  loss  of  iodine  by  the  formation  of  iodate.  It  excludes 
any  explanation  of  the  red  color  based  upon  the  assumed  forma- 
tion of  some  arsenic  acid  compound  with  starch. 

It  is  well  known  that  malt  extract;  and  many  chemical  re- 
agents such  as  hydrochloric  acid,  potassium  hydroxide,  nitric 
acid,  etc.,  readily  hydrolyze  starch  more  or  less  completely 
through  a  series  of  bodies,  called  dextrins,  to  a  final  product, 
one  of  the  sugars.  One  of  the  first  of  these  dextrins  is  erythro- 
dextrin,  which  is  colored  red  with  iodine.  It  seems  plausible 
that  water  or  the  alkali,  acid  potassium  carbonate,  might  cause 
hydrolysis  under  certain  conditions.  As  there  is  a  loss  of  iodine, 
indications  would  point  to  an  oxidizing  action,  and  a  hydrolysis 
may  be  possible  because  of  such  oxidizing  action.  It  is  also 
possible  that  the  arsenious  acid  (or  aritimonious  acid)  becomes 
joined  to  the  starch,  in  some  such  way  as  antimonious  acid 
attaches  to  acid  potassium  tartrate,  and  that  this  compound  is 
then  easily  hydrolyzed.  A  possible  indication  of  this  lies  in  the 
anomalous  fact  that  the  blue  color  fades  first  and  the  red  last 
in  titration  in  alkaline  solution,  though  a  few  drops  of  a  starch 
solution  added  to  a  solution  of  erythrodextrin,  colored  red  with 
iodine,  will  develop  the  blue  starch  iodide. 

Experimental  evidence  seems  to  substantiate  the  following 
points : 

1st.  The  loss  of  iodine  and  the  production  of  a  red  color  does 
not  take  place  if  an  absolutely  pure  and  freshly  made  starch 
solution  is  employed. 

2d.  Ordinary  starch  contains,  usually,  at  least  two  impurities, 
one  coloring  red  with  iodine,  the  other  coloring  blue,  the  latter 
being  readily  changed  under  the  influence  of  oxygen  and  acid 
potassium  carbonate  to  the  former.  These  impurities  tend  like- 
wise to  form  in  pure  starch,  whether  solid  or  in  solution. 

3d.  The  impurity  coloring  blue  with  iodine  is  identical  with, 
or  analogous  to,  amidulin,  made  by  saliva  digestion  of  pure 
starch,  and  the  impurity  coloring  red  with  iodine  is  erythro- 
dextrin, the  second  product  of  saliva  digestion  of  pure  starch. 


32  METHODS  IN  CHEMICAL  ANALYSIS 

4th.  The  loss  of  iodine  is  due  to  the  formation  of  erythro- 
dextrin  from  this  amidulin-like  body,  and  erythrodextrin  does 
not  use  up  iodine  by  any  transformation  to  achroodextrins. 

So  it  appears  that  the  colors  found  in  iodometric  titrations  in 
which  ordinary  starch  is  used  as  an  indicator  are  probably  due 
to  the  admixtures  of  the  starch  blue  or  possibly  of  the  amidulin 
blue  with  the  red  of  erythrodextrin  derived  from  amidulin  by 
hydrolysis  initiated  by  the  oxidizing  effects  of  the  iodine.  Pure 
starch,  containing  neither  amidulin  nor  erythrodextrin,  gives  only 
blue  in  the  iodometric  titration.  Starch,  on  the  other  hand, 
which  has  undergone  partial  hydrolysis,  is  likely  to  contain  both 
amidulin  and  erythrodextrin.  It  is  not  strange  that  both  amidu- 
lin and  erythrodextrin  should  be  present  as  impurity  in  starch  > 
since  they  both  stand  in  the  order  named  as  the  first  two  dextrins 
produced  from  starch,  as  shown  by  saliva  digestion  of  starch,  as 
also  by  malt-extract  digestion  of  starch.  Starch  both  in  solid 
state  and  in  solution  tends  to  pass  through  these  stages  of  hy- 
drolysis. Germ  growth  rapidly  appears  in  solutions  of  pure 
amidulin  and  pure  erythrodextrin,  with  the  destruction  of  these 
bodies  to  form  dextrins  lower  in  the  series. 

Pure  starch  causes  no  red  color,  nor  loss  of  iodine,  in  alkaline 
titration  of  arsenite  solution  or  of  tartar  emetic.  If  any  purplish 
tinge  occasionally  occurs,  it  is  no  hindrance  to  the  reading  and 
causes  no  appreciable  loss  of  iodine. 

With  an  impure  starch,  the  reading  from  the  first  permanent 
color,  whether  red  or  blue,  is  nearest  to  the  correct  value.  The 
readings  may  be  compared  with  plain  iodine  readings  and  a  cor- 
rection applied,  since  the  loss  for  a  constant  quantity  of  starch 
proves  to  be  constant  in  the  titration  of  20  cm.3  to  50  cm.3 
of  arsenite  solution.  Titration  should  be  made  at  considerable 
dilution  —  e.g.,  150  cm.3  to  200  cm.3  (since  the  production  of  red 
is  at  a  minimum  and  the  loss  of  iodine  small  at  high  dilutions)  — 
and  in  presence  of  a  suitable  amount  of  potassium  iodide  to  ren- 
der the  reaction  delicate.  Whenever  it  is  practicable,  however, 
the  best  method  of  using  an  impure  starch  is  to  make  the  titra- 
tion without  it  up  to  the  appearance  of  the  yellow  tinge  of  free 
iodine  and  then  to  add  the  starch.  In  this  way  the  color  comes 
out  a  clear  blue  and  the  exact  adjustment  may  be  easily  made 
by  alternate  additions  of  a  drop  or  two  of  n/io  arsenite  and  w/io 
iodine. 


APPLIANCES  AND   GENERAL  PROCEDURE  33 

Hale  *  experimented  with  various  preparations  of 
Ordinary  Prepa-  starch.  The  starch  solution  was  made  by  grinding 
5  grm.  of  starch  paste  with  a  few  cubic  centimeters 
of  cold  water  with  the  addition  of  o.oi  grm.  of  mercuric  iodide, 
pouring  into  a  liter  of  boiling  water,  and  boiling  five  to  ten  min- 
utes. Only  the  clear  supernatant  liquid  was  used.f 
Preparation  in  25  cm.3  of  cold  water  5  grm.  of  pure  starch 

with  Potassium  .   ,  ..  .....  , 

iodide.  were  ground  with  2  grm.  of  potassium  iodide  and 

the  mixture  was  poured  into  75  cm.3  of  boiling  water  and 
boiled,  the  beaker  being  protected  with  asbestos.  The  mass 
became  mucilaginous.  After  fifteen  minutes  the  volume  was  in- 
creased to  500  cm.3,  and  the  boiling  was  continued  for  forty-five 
minutes.  The  solution  was  filtered,  forty-eight  hours  being  re- 
quired, leaving  a  residue  of  jellylike  consistency  upon  the  filter. 
This  method  was  suggested  by  the  extreme  delicacy  which  the 
presence  of  potassium  iodide  gives  to  the  starch  reaction,  and  by 
certain  statements  made  in  the  literature,  one  that  concentrated 
potassium  iodide  causes  starch  to  swell  and  dissolve,!  another 
that  a  solution  of  starch  made  by  a  somewhat  similar  method 
would  keep  a  year  without  fermentation. § 

Heating  with  In  70  cm.3  of  pure  glycerin  5  grm.  of  potato  starch 
ForT^o-  were  heated  at  a  temperature  of  i85°-i9O°  C.  for  half 
dextrin.  an  hour  with  constant  stirring.  The  starch  dissolved 

and  the  solution  turned  through  yellow  to  a  deep  red.  The  solu- 
tion was  cooled  to  120°  C.  and  poured  slowly  and  continuously 
into  200  cm.3  of  alcohol.  The  precipitate  was  thoroughly  stirred, 
settled,  and  filtered  while  warm.  It  filtered  readily  and  was 
washed  with  alcohol  until  the  filtrate  came  through  colorless. 
The  colorless  residue  of  amorphous  amylodextrin  was  then  dis- 
solved in  500  cm.3  of  water  heated  to  6o°-7O°  C. 

Soluble  starch  *n  a  ^tt:^e  co^  water  2  grm-  °f  starch  and  0.5  grm. 
by  saliva  Diges-  of  acid  potassium  carbonate  were  ground  together 
tion  (Amiduiin).  an£  mixed  with  200  cm.3  of  water  kept  boiling  for  a 
few  minutes.  The  mixture  was  cooled  to  4O°-45°  and  treated 
with  10  cm.3  of  filtered  saliva  neutralized  byo.i  percent  hydro- 
chloric acid  with  the  use  of  a  slip  of  litmus  paper  (first  dipped  in 

*  Loc.  cit. 

t  Gastine's  formula,  Zeit.  anal.  Chem.,  xxviii,  339. 

J  Payen,  Compt.  rend.,  Ixi,  512. 

§  Zeit.  anal.  Chem.,  xxv,  37. 


34 


METHODS  IN  CHEMICAL  ANALYSIS 


acetic  acid  and  then  washed)  as  indicator.  When,  in  the  course 
of  three  or  four  minutes,  the  solution  had  become  entirely  clear, 
it  was  boiled  for  ten  minutes.  In  this  process  the  addition  of 
the  alkali  hinders  the  action  from  going  beyond  the  first  step 
of  digestion:  the  boiling  at  the  end  destroys  any  further  action 
of  the  saliva.  Starch  cellulose,  which  is  said  to  produce  a  feeble 
red  or  brownish  color  with  iodine,*  is  digested  and  destroyed  by 
the  saliva. 

In  the  presence  of  a  suitable  amount  of  potassium  iodide  each 
of  the  first  three  preparations  gives  a  sharp  indication  with  a 
single  drop  of  n/io  iodine;  the  amidulin  is  only  a  little  less  deli- 
cate. In  the  titration  of  50  cm.3  of  n/io  arsenite,  enough  potas- 
sium iodide  is  present  in  the  n/io  iodine  to  give  a  sharp  reaction 
at  a  volume  of  125  cm.3  This  is  shown  in  the  following  state- 
ment. For  comparison,  a  titration  made  with  a  large  amount 
(25  cm.3)  of  ordinary  impure  starch  is  included. 

Titrations  with  the  Different  Preparations  of  Starch. 
Volume  i25Tcm.3. 


Starch  solution. 

cm.» 

W/IO 

Nearly 

^o!0i 

w/io  I  sol. 

Pure 

Impure 

Color. 

KI 

ordi- 

Amylo- 

Amid- 

ordi- 

starch. 

nary 

dextrin. 

ulin. 

nary 

starch. 

starch. 

cm.1 

cm.s 

5° 

49.38 

I 

Permanent  purplish. 

5° 

49.40 

I 

( 

Good  blue. 

5° 

49.40 

I 

Good  blue. 

50 

49.40 

2 

Slow-fading  purplish. 

5° 

49.42 

. 

2 

Good  blue. 

5° 

.  . 

1-5 

Permanent  purplish. 

So 

49.40 

1-5 

Deep  blue. 

So 

49.42 

5 

Permanent  purplish. 

So 

49-44 

. 

5 

Purple. 

50 

49-55 

25 

Abundant  red. 

! 

In  a  series  of  parallel  titrations,  designated  A  and  B  in  the 
next  table,  readings  were  made  with  the  blue  developed  by  the 
Kl-starch  indicator  and  with  plain  iodine,  alternately,  in  order 
to  eliminate  accidental  errors.  Starch  was  added  subsequently 
to  corroborate  the  very  pale  iodine  reading.  The  corrected 
readings  were  found  by  subtracting  one  drop  from  the  actual 

*  Beilstein  (First  Edition),  i,  1082,  line  17. 


APPLIANCES  AND    GENERAL  PROCEDURE 


35 


reading.     To  render  the  plain  iodine  readings  sharp  a  crystal  of 
potassium  iodide  was  added  in  the  first  two  titrations. 


Titrations  -with  and  without  Starch. 
Volume  125  cm.3 


Nearly  M/IO  I. 
Direct  readings. 

Nearly  w/io  I. 
Corrected  readings. 

M/IO  I  sol. 
Absolute 

M/IO  I  SOl. 

Absolute  errors 
in  A,  .  .  in  B. 

H/IO 

As203. 

By  iodine 
color, 
A. 

ByKI- 
starch 
blue, 
B. 

A. 

B. 

amount 
calculated 
from 
50  cm.3 

A. 

B. 

cm.8 

cm.3 

cm.s 

cm.3 

cm.3 

cm.3 

cm.8 

cm.3 

i  drop* 

i  drop 

5 

4-94* 

4.96 

4.92 

4-94 

4-94 

—  O.O2 

o.oo 

10 

9.88 

9.90 

9.86 

9.88 

9.88 

—  O.O2 

o.oo 

15 

14-83 

14-83 

14.81 

14.81 

14.82 

—  O.OI 

—  O.OI 

20 

19.78 

19.78 

19.76 

19.76 

19.76 

0.00 

o.oo 

30 

29.64 

29.64 

29.62 

29.62 

29-63 

—O.OI 

—  O.OI 

40 

39-55 

39-55 

39-53 

39-53 

39-51 

+  0.02 

+O.O2 

So 

49.41 

49.41 

49-39 

49-39 

49-39 

o.oo 

0.00 

*  A  crystal  of  KI  was  added. 

It  is  to  be  noted  that  the  readings  with  plain  iodine  and  with 
pure  starch  agree  exactly  except  for  the  first  two  titrations,  and 
here  there  is  only  a  difference  of  a  drop.  The  absolute  errors  are 
interesting,  as  they  show  how  the  absolute  values  fluctuate  about 
a  standard  set  by  the  5O-cm.3  readings.  This  fluctuation  is  lim- 
ited to  a  drop  plus  or  minus. 

The  statement  has  been  made  that  starch  from  different  sources 
has  a  varying  power  of  absorbing  iodine,  e.g.,  that  potato  starch 
absorbs  three  times  as  much  as  rice  starch.*  To  discover 
whether  this  fact  has  any  bearing  upon  the  use  of  the  starch  in- 
dicator, and  at  the  same  time  to  learn  whether  pure  starch  solu- 
tions made  in  the  ordinary  way  (Gastine)f  would  give  as  delicate 
readings  as  that  boiled  with  potassium  iodide,  solutions  were 
made  from  pure  potato  starch,  pure  rice  starch,  pure  arrowroot 
starch  and  a  pure  (so-called)  soluble  starch  of  unknown  origin. 
The  results  of  titration  with  these  solutions  as  indicators  are 
shown  in  the  following  statement. 

It  is  at  once  seen  that  these  values  are  coincident  within  a 
drop,  and  that  all  the  starch  solutions  are  within  the  limits  set 
by  a  plain  iodine  reading  on  the  one  hand  and  a  potassium  iodide 

*  Girard,  Ann.  Chim.,  [6],  xiii,  275. 
t  Loc.  cit. 


METHODS  IN    CHEMICAL  ANALYSIS 


starch  reading  on  the  other,  though  there  was  no  dilution  of  the 
standard  solutions. 

Starch  from  Various  Sources. 

No  extra  dilution.  Volume  about  no  cm.8 


W/IO 

As20,. 
cm.* 

Nearly 
«/io  I  sol. 

cm.8 

KHCO3 

satu- 
rated 
cm.8 

Starch  solution. 
cm.8 

Color. 

50 

50 
5° 
SO 

50 
CQ 

{49."  I 
I  49  •  13  ) 
49-13 
49.14 
49-15 

49-15 
40    1  3 

5 

5 
5 
5 

5 

e 

i  (Ordinary  pure  potato.)  < 

i  (Ordinary  pure  rice.) 
i  (Ordinary  pure  soluble.) 
i  (Ordinary  pure  arrow- 
root.) 
i  (KI  pure  starch.) 

Pale  blue. 
Deep  blue. 
Blue,  slightly  purplish. 
Blue,  slightly  purplish. 

Deep  blue. 
Good  blue. 
Yellow 

If  the  starch  is  pure  the  amount  of  it  used,  within  reasonable 
limits,  is  without  effect  upon  the  titration  of  n/io  arsenite  by 
n/io  iodine,  and  the  same  may  be  said  of  the  amount  of  acid 
carbonate,  as  shown  in  the  following  table. 

Variation  in  the  Amount  of  Starch. 


M/IO  AsjOa. 
cm.J 

Nearly  «/  10  1  sol. 
cm.8 

KHCO3 

saturated. 
cm.8 

Starch  solution.* 
cm.8 

Color. 

10 

9.82 

5 

i-5 

Deep  blue,  purplish. 

IO 

9.82 

5 

5 

Deep  blue. 

10 

jp.Bal 
19.24) 

5 

•°    1 

Pale  blue,  purplish. 
Deep  blue. 

10 

9.82 

5 

15 

Deep  blue. 

IO 

9.84 

5 

20 

Deep  blue. 

10 

9-85 

5 

25 

Deep  blue. 

10 

9.82 

5 

•  5 

Deep  blue. 

10 

9.82 

10 

•  5 

Deep  blue. 

10 

9-83 

15 

-5 

Deep  blue. 

IO 

9.82 

20 

•  5 

Deep  blue,  purplish. 

10 

9.81 

25 

•  5 

Deep  blue. 

*  Ordinary  preparation;  pure  potato  starch. 

With  pure  starch,  the  end  reaction  of  the  titration  of  tartar 
emetic  by  iodine  is  likewise  a  pure  blue.  Following  are  Hale's 
results. 

The  average  of  the  10  cm.3  readings  (absolute)  multiplied  by 
five  equals  47.73.  The  absolute  5O-cm.3  reading  (47.75-0.02) 
equals  47.73.  Evidently  even  tartar  emetic  causes  no  loss  on 
pure  starch,  for  the  5O-cm.3  reading  agrees  with  the  plain  iodine 
reading  for  the  same  amount  and  with  the  io-cm.3  titrations. 


APPLIANCES  AND   GENERAL  PROCEDURE 


37 


End  Reaction  with  Tartar  Emetic. 


Volume. 

«/io  tartar 
emetic. 

Nearly 
w/io  I  sol. 

Starch 

solution, 

KHCO3. 

Color. 

pure  potato. 

cm.3 

cm.1 

cm.* 

cm.8 

cm.» 

100 

10 

9.58 

i-5 

10 

Blue,  no  red. 

IOO 

10 

9.58 

i-5 

10 

Blue. 

75 

IO 

9-S6 

IO 

Yellow. 

125 

50 

47-75 

i-5 

25 

Blue  (purplish  tinge). 

125 

50 

47-75 

25 

Yellow. 

Hale*  emphasizes  the  need  of  potassium  iodide  in  suitable 
amount  to  bring  out  the  delicacy  of  the  blue  end  reaction,  and 
the  further  necessity  of  not  exceeding  a  suitable  proportion,  in 
order  that  the  starch  blue  compound  may  not  be  modified  by 
transformation  to  starch  red.  The  proportions  of  potassium 
iodide  and  iodine  entering  into  starch  blue  and  starch  red  were 
carefully  studied.  According  to  experimental  results,  it  would 
appear  that  the  group  KI.I4  is  characteristic  of  starch  blue  and 
the  group  KI.I2  of  starch  red ;  that  in  the  presence  of  a  sufficiently 
concentrated  solution  of  potassium  iodide  the  group  KI.I4  changes 
by  addition  of  KI  to  2KI.I2,  while  dilution  with  water  tends  to 
bring  about  the  reverse  action. 

Wholly  apart  from  the  consideration  of  theory,  the  influence 
of  iodides  upon  the  delicacy  of  the  starch  blue  test  for  iodine  is 
a  matter  of  considerable  practical  importance  analytically.  Ex- 
perience shows  that  0.3  grm.  of  potassium  iodide  is  sufficient  in 
volumes  not  exceeding  300  cm.3  to  make  readings  by  the  starch 
indicator  sharp,  and  that  under  these  conditions  indications  are 
quite  as  delicate  at  ordinary  room  temperatures  as  at  the  tem- 
perature of  ice  water.  The  addition  of  more  potassium  iodide 
renders  the  reading  no  sharper  within  that  range  of  dilution, 
though  the  total  volume  has  considerable  proportionate  influence 
upon  the  amount  of  iodine  needed  to  bring  out  the  indication. 
In  the  absence  of  potassium  iodide  in  sufficient  amount,  the  in- 
fluence of  temperature  is  very  noticeable.  Hydrochloric  acid 
either  in  small  amount  or  large  amount  does  not  render  the  read- 
ing sharp  in  absence  of  potassium  iodide.  These  results  are  in 
agreement  with  the  work  of  Lonnes.f 


*  Am.  Chem.  Jour.,  xxviii,  450. 


t  Zeit.  anal.  Chem.,  33,  409- 


METHODS  IN  CHEMICAL  ANALYSIS 


Effect  of  Temperature  when  Potassium  Iodide  is  Restricted. 


Volume. 

Potassium 
iodide  not 
exceeding 

Hydro- 
chloride 
acid. 
(Sp.gr.  1.  12.) 

Pure 

potato 
starch 
solution. 

Temper- 
ature. 

M/IOOO 

iodine. 

Color. 

0.3  grm. 

cm.* 

cm.3 

cm.* 

cm.* 

100 

i  crystal 

2 

23° 

0-45 

Faint  blue. 

200 

i  crystal 

2 

23° 

0.65 

Faint  blue. 

300 

i  crystal 

2 

23° 

0.80 

Faint  blue. 

400 
300 

i  crystal 
i  crystal 

2 

2 

23° 

23° 

0.80 

Faint  blue. 
Faint  blue. 

300 

i  crystal 

2 

5° 

0.70 

Faint  blue. 

300 

2 

23° 

9.  20 

Faint  blue. 

300 

2 

5° 

I    'CO 

Faint  blue 

300 

I 

2 

23° 

A    ie 

Faint  blue 

300 

IO 

2 

23° 

4OO 

Faint  blue 

The  influence  of  different  amounts  of  potassium  iodide  in 
bringing  out  the  starch  iodide  indication  of  iodine  set  free  by 
reaction  with  gold  chloride  is  of  interest  in  this  connection.  At 
volumes  lying  between  the  limits  of  25  cm.3  and  50  cm.3  o.i  grm. 
of  potassium  iodide  is  an  appropriate  amount;  at  a  volume  of 
15  cm.3,  o.oi  grm.  to  0.05  grm.  of  the  iodide  will  do  the  work; 
and  at  lower  dilutions  even  less  of  the  iodide  is  effective.* 


Variation  of 
Standard. 


Standard  Tartar  Emetic. 

Gruenerf  has  shov/n  that  tartar  emetic  solutions 
containing  about  1 6  grm.  of  tartar  emetic,  20  grm. 
to  30  grm.  of  tartaric  acid  and  I  cm.3  of  concentrated  hydro- 
chloric acid  to  the  liter,  will  keep  from  five  to  twelve  months 
without  any  change  in  strength.  There  is  no  deposit  of  anti- 
monious  oxide  under  these  conditions,  no  oxidation  and  no  signs 
of  fungous  growth.  Gruener  determined  the  strength  of  his 
tartar  emetic  solutions  by  titration  with  a  decinormal  iodine 
solution,  standardized  by  decinormal  arsenite.  The  mean  of 
twenty-nine  determinations  showed  43.95  per  cent  of  antimo- 
nious  oxide  in  tartar  emetic.  Theory  required  43.37  per  cent 
(Sb  =  i20,  KSbOC4H4O64H2O  =  332).  The  cause  of  this  dis- 
crepancy between  arsenite  and  tartar  emetic  solutions  made  up 
as  standards  according  to  the  accepted  molecular  formulas  is  a 
matter  of  considerable  interest.  One  suggested  explanation  of 

*  See  page  146. 

t  Hippolyte  Gruener,  Am.  Jour.  Sci.,  [3],  xlvi,  206. 


APPLIANCES  AND   GENERAL  PROCEDURE  39 

this  difference  is  that  the  end  reaction  between  starch  and  iodine 
is  delayed  until  an  excess  of  iodine  is  present.  Hale*  has  shown, 
however,  that  a  pure  starch  solution  gives  a  sharp  end  reaction 
with  both  tartar  emetic  and  arsenite  solutions,  and  that  while 
with  impure  starch  there  is  a  loss  of  iodine  accompanied  by  the 
production  of  reddish  hues  in  titrating  tartar  emetic,  as  shown 
by  the  difference  between  the  readings  made  in  the  presence  of 
potassium  iodide  by  the  yellow  color  of  iodine  and  by  the  blue 
of  starch  iodide,  yet  it  is  no  greater  than  in  titrating  arsenite 
solution  in  the  presence  of  an  impure  starch.  If  only  an  impure 
starch  is  available,  the  reading  should  be  made  without  starch, 
for  the  presence  of  potassium  iodide  renders  very  sharp  the  yel- 
low color  of  the  first  excess  of  free  iodine.  This  first  reading 
may  be  afterwards  corroborated  by  adding  the  starch  solution, 
which  will  then  give  only  a  pure  blue  color.  The  above  dis- 
crepancy must  then  be  due  to  some  other  cause  than  delay  of 
formation  of  the  starch  iodide.  Halef  has  given  proof  that  this 
discrepancy  is  due  to  the  ease  with  which  tartar  emetic  loses  its 
water  of  crystallization,  and  that  in  order  to  get  a  salt  of  the 
exact  composition,  KSbOC4H4O6.^H2O  (mol.  wt.  332.15),  certain 
conditions  must  be  very  closely  observed. 

Molecular  weights  of  tartar  emetic  variously  prepared,  calcu- 
lated from  the  results  of  iodometric  titration,  run  in  a  series  from 
that  of  crystalline  tartar  emetic  almost  to  that  the  salt  which 
has  lost  i  .5  molecules  of  water,  passing  through  all  intermediate 
stages,  but  never  surely  resting  at  any  one  spot.  Two  important 
stages  are  reached :  when  all  the  water  of  crystallization  is  gone, 
the  anhydrous  state ;  and  when  0.5  molecule  of  water  further  has 
been  lost,  the  first  anhydride  stage.  The  variation  of  molecular 
weight  is  shown  in  the  following  table. 

The  condition  most  easily  and  definitely  reached  is  that  of 
the  hydrous  crystalline  salt.  The  greatest  error  met  with  in 
the  recrystallized  salt,  if  air-dried,  is  about  +0.2  per  cent,  calcu- 
lated on  the  ratio  of  antimony  to  tartar  emetic,  and  that  after 
standing  in  fine  condition  in  a  closed  bottle  for  several  weeks. 
Drying  in  the  air  bath  does  not  yield  a  product,  even  with  the 
most  finely  divided  preparation, — that  precipitated  by  alcohol, — 
which  is  sufficiently  uniform  to  serve  the  purpose  of  a  standard. 

*  F.  E.  Hale,  Am.  Jour.  Sci.,  [4],  xiii,  379. 

t  F.  E.  Hale,  Jour.  Am.  Chem.  Soc.,  xxiv,  828. 


METHODS  IN  CHEMICAL  ANALYSIS 


Hale  recommends,  therefore,  the  preparation  of  tartar  emetic, 
in  medium-sized  crystals  (fa  to  J  inch  in  diameter) ,  rather  than 
in  the  minutely  crystallized  condition. 

Variations  in  Molecular  Weight. 

49.29  cm.3  (absolute  amount)  of  w/io  iodine  solution  =  50  cm.3  of  w/io  ar- 
senic trioxide  solution. 


Preparation. 

n/io 
tartar 
emetic 
solu- 
tion. 

cm.» 

W/IO 

iodine 
solu- 
tion. 

cm.3 

Molec- 
ular 
weight. 

Anti- 
mony. 

per  cent. 

Remarks. 

Freshly  crystallized  

CO 

49.  29 

6  i 

' 

O 

49.56 

330.34 

Crystals  kept  several  weeks  .... 

(  5o 

49.58 

330.21 

36.35 

Tartar 

emetic. 

grm. 

Medium-sized  crystals,  air- 

dried                                         .  . 

0.5 

20    82 

"2  3O     -  S 

Crystals,  4  days  over  ^SCV.  .  .  . 

0-5 

^v  •  u 

29.95 

329.13 

Crystals   7  days  over  H2SO4. 

o  •  5 

20    08 

328   77 

Crystals,  16  days  over  H2SO4. 

o  •  5 

~y  *  y 
•JQ  16 

o      •  /  / 
327   l< 

Crystals,  4  hours  at  95°-i3o°.  .  . 
Crystals,  7?  hours  at  io4°-i3o°.. 

0-5 

0-5 

ow* 
30.42 

30.51 

O      t         3 
324.50 
323.I5 

37-14 

(Anhydrous.) 

Crystals                  at  I04°-i3o° 

j  °-5 

31.27 

3I5.29 

(  ... 

314.  15 

(£H2O  anhy- 

dride.) 

Crystals.                at  i6o°-i65° 

0.5 

31.88 

309  .  28 

Crystals,  2  hours  at  i6o°-i65°.  . 

3L-93 

308.80 
305-I5 

(H2O  anhy- 

dride —  all 

hydroxyls 

gone.) 

Preparation. 


Enough  tartar  emetic  is  dissolved  in  about  300  cm.3 
of  boiling  water  to  make  a  concentrated  but  not  a 
saturated  solution.  This  hot  solution  is  filtered  into  flat  crystal- 
lizing dishes  and  allowed  to  crystallize  over  night.  The  crystal- 
lization should  not  be  too  rapid.  The  crystals  are  filtered  off  by 
suction,  washed  twice  with  distilled  water,  kept  under  suction 
for  about  five  or  ten  minutes  more,  and  then  air-dried  from  one 
and  a  half  to  four  hours  at  room  temperature,  not  above  25°  and 
preferably  lower,  in  a  clear,  dry  air.  The  crystals  may  be  con- 
sidered dry  an  hour  or  two  after  they  cease  to  show  the  slightest 
tendency  to  cling  to  a  glass  rod  used  as  a  stirrer. 

Tartar  emetic  may  be  prepared  in  this  manner,  which,  with 
good  starch,  shows  practically  the  theoretical  value  when  tested 


APPLIANCES  AND   GENERAL  PROCEDURE 


against  a  standard  arsenite  solution  by  titration  of  both  with 
iodine.  As  has  been  pointed  out,  if  pure  starch  is  not  available 
the  first  reading  should  be  made  without  starch  in  presence  of 
potassium  iodide,  and  this  corroborated  after  adding  the  starch 
solution,  which  will  then  give  only  a  pure  blue. 

A  comparison  of  several  preparations  of  tartar  emetic  with 
standard  arsenite  is  given  in  the  table. 

Comparison  of  Tartar  Emetic  with  Standard  Arsenite. 


.2 

I 

^ 

ll 

n 

Crystals,  air- 

i*3 

'«  o 

jj  o> 
c5  a 

a  J2 

"S 

w/io  solution. 

dried. 

la 

Rl    O 

8-3  d 
'S.2.2 

.2  a) 

1 

Color. 

SP 

*J^ 

Jl 

u 

^S 

K  12 

5*  H    OT 

(2*° 

55 

hours,    temp. 

cm.3 

cm.3 

cm.3 

cm.3 

cm.3 

Tartar  emetic  1  

ii 

I90-24° 

5Q 

49-32 

2S 

Medium  blue. 

Tartar  emetic  I  

4 

i9°-24° 

5° 

49.42 

25 

Medium  blue. 

Tartar  emetic  II..  . 
Arsenite 

i9°-24° 

5° 

49-43 
49  43 

25 

5" 

Medium  blue. 
Medium  blue. 

Tartar  emetic  III.  . 
Tartar  emetic  IV  .  . 

3 
4 

i9°-24° 

5o 

49.58 
49.58 

25 
25 

Deep  blue. 
Medium  blue. 

Arsenite  .         ... 

50 

49.61 

5 

Medium  blue. 

PROCESSES  OF  OXIDATION. 
Arsenic  Trioxide  as  a  Standard. 

Carefully  sublimed  and  anhydrous  arsenic  trioxide  serves  admi- 
rably for  the  standardization  of  potassium  permanganate  for  use 
in  direct  oxidations,  as  well  as  in  the  processes  of  iodometric 
analysis.  It  is  generally  best  to  work  with  standard  arsenite 
solution  prepared  as  previously  described,*  and  it  is  convenient, 
though  not  necessary,  to  have  at  hand  a  solution  of  iodine, f  also 
standardized  against  the  arsenite  solution. 

standardization  The  standardization  of  the  permanganate  solu- 
without  iodine,  tion,  made  up  to  an  approximate  value,  is  easily 
accomplished  by  running  a  suitable  amount  of  it  into  a  solution 
of  potassium  iodide  contained  in  a  reaction  bottle  {  and  acidi- 
fied with  dilute  sulphuric  acid,  neutralizing  with  acid  potassium 

*  See  page  29. 

f  Ibid. 

I  See  page  6. 


42  METHODS  IN  CHEMICAL  ANALYSIS 

carbonate,  and  titrating  by  the  standard  arsenite  the  iodine  set 
free  by  the  permanganate.  In  this  operation  it  is  best  to  dis- 
pense with  the  starch  indicator  usually  employed  to  fix  the  end 
reaction.  The  vanishing  point  of  the  color  of  free  iodine  is  itself 
sufficiently  definite,  even  at  a  dilution  of  300  cm.3,  and  the  dis- 
appearance of  color  is  much  sharper  than  that  of  the  blue  starch 
iodide.* 

standardization  ^  a  solution  of  iodine  standardized  against  the 
with  the  Aid  of  arsenite  solution  is  at  hand,  the  process  just  de- 
scribed may  be  modified  so  that  there  is  no  danger 
of  overrunning  the  end-point  in  a  single  titration.  In  this  pro- 
cedure, an  excess  of  the  standard  arsenite,  taken  in  known 
amount,  is  added  at  once  after  the  reaction  of  the  permangan- 
ate upon  the  iodide  in  presence  of  acid,  and  is  followed  by  the 
acid  carbonate.  The  excess  of  arsenite  is  determined  by  titration 
with  iodine  in  presence  of  starch.  Should  too  much  iodine  be 
added  in  the  titration,  it  is,  of  course,  only  necessary  to  add 
another  measured  amount  of  arsenite  and  then  to  repeat  the 
titration  by  iodine. f 

The  Gravimetric  Standardization  of  Permanganate. 

The  standardization  of  a  permanganate  solution  may  be  made 
gravimetrically  by  adding  a  suitable  amount  of  the  solution  to 
an  excess  of  potassium  iodide  acidified  with  hydrochloric  acid, 
shaking  the  mixture  with  a  weighed  amount  of  specially  prepared 
silver  in  an  atmosphere  of  hydrogen,  collecting  upon  asbestos 
in  the  perforated  crucible  the  residue  of  silver  and  silver  iodide, 
drying,  and  weighing,  according  to  the  procedure  previously 
described.}: 

The  Loss  of  Oxygen  in  Oxidations  by  Potassium  Permanganate. 

Concentration  A  statement  made  many  years  ago,  that  in  the 
of  Acid.  interaction  of  oxalic  acid  and  potassium  perman- 

ganate free  oxygen  is  always  a  product, §  met  with  adverse  criti- 
cism ;  but  subsequently  a  similar  effect  was  noted  by  Brauner  11 

*  Gooch  and  Peters,  Am.  Jour.  Sci.,  [4],  viii,  125. 
f  Gooch  and  Gilbert,  Am.  Jour.  Sci.,  [4],  xv,  390. 
j  See  page  27. 

§  Francis  Jones,  Jour.  Chem.  Soc.,  1878,  95. 
||  Jour.  Chem.  Soc.  (1891),  238. 


APPLIANCES   AND    GENERAL   PROCEDURE 


43 


in  the  action  of  permanganate  upon  tellurous  acid  dissolved  in 
sulphuric  acid,  with  the  additional  observation  that  the  evolu- 
tion of  oxygen  is  proportional  to  the  amount  of  sulphuric  acid 
employed,  and  that  in  alkaline  solution  little  evidence  of  such  an 
effect  appears.  Recognizing  that  the  production  of  permanganic 
acid,  free  oxygen  and  ozone,  by  the  action  of  strong  sulphuric 
acid  upon  permanganate  in  absence  of  oxidizable  material,  is  a 
common  phenomenon,  and  that  the  formation  of  a  precipitate 
consisting  largely  of  hydrated  manganese  dioxide  by  the  action 
of  hot  dilute  sulphuric  acid  upon  the  permanganate  in  aqueous- 
solution  is  likewise  well  known,  Gooch  and  Danner*  have  in- 
vestigated the  action  of  sulphuric  acid  in  different  concentrations 
upon  permanganate  in  solution  with  a  view  to  determining  how 
far  such  action  may  be  directly  or  indirectly  responsible  for  the 
liberation  of  free  oxygen  in  processes  of  oxidation. 

In  certain  experiments  tubes  of  suitable  size  and  length,  hold- 
ing from  100  cm.3  to  200  cm.3,  were  sealed  at  one  end,  filled  com- 
pletely with  the  mixtures  of  acid  and  permanganate,  inverted > 
and  allowed  to  stand  with  the  lower  and  open  end  submerged 
in  liquid  of  the  exact  composition  of  that  which  filled  them.  The 
details  and  results  of  these  experiments  are  recorded  in  the  state- 
ment below. 


Time 
elapsed. 

Gas  from 
100  cm.J 

Appearance. 

Time 
elapsed. 

Gas  from 
ico  cm.3 

Appearance. 

A. 

B. 

H2S04  [i  :  i]=5o  per  cent. 

H2SO4  [i  :  i]=25  per  cent. 

S  rain. 

o.i  cm.1 

No  change. 

S  min. 

Small  bubble. 

No  change.     """ 

I  hour. 

i.i  cm.* 

No  change. 

i  day. 

14     cm.1 

Red  brown. 

•  •  ,  •  ' 

3  days. 
4  days. 

15.  3  cm.* 
15.  6  cm.* 

Light  brown. 
Light  brown. 

3  days. 

9.6  cm.8     | 

Reddish  purple. 
Turbid. 

7  days. 

Brown,  turbid. 

( 

Reddish  pink. 

8  days. 

16     cm.* 

(Clearing  by  precipi- 
\     tation. 

7  days. 
15  days. 

15.  i  cm.*     < 
18     cm.* 

Clearing  by  pre- 
cipitation. 
Nearly  clear. 

15  days. 

17.  3  cm.* 

Clear,  straw-colored. 

17  days. 
35  days. 

17.4  cm.* 
17.5  cm.* 

Clear,  straw-colored. 
Clear,  straw:colored. 

35  days. 

18.4  cm.3     | 

Clear  and  color- 
less. 

C. 

D. 

H2SO4  [i  :  i]  =  i2.s  per  cent. 

H2SO4  (i  :  1  1=6.25  per  cent. 

i  hour. 

Small  bubble. 

No  change. 

i  hour. 

Small  bubble. 

No  change. 

i  day. 
3  days. 
14  days. 
37  days. 
44  days. 

Bubble. 
Bubble  larger. 
7.1    cm.1 
ii    cm.* 
12    cm.* 

No  change. 
No  change. 
Color  lighter. 
Color  lighter. 
Color  lighter. 

i  day. 
3  days. 
14  days. 
37  days. 
44  days. 

Bubble. 
Bubble  larger, 
i  cm.3 
3  cm.3 
5  cm.* 

No  change. 
No  change. 
No  change. 
No  change. 
Little  change. 

*  F.  A.  Gooch  and  E.  W.  Danner,  Am.  Jour.  Sci.,  [3],  xliv,  301. 


44 


METHODS  IN   CHEMICAL  ANALYSIS 


In  other  experiments  note  was  made  of  changes  in  color  and 
formation  of  precipitates  in  ioo-cm.3  portions  of  liquid  containing 
10  cm.3  of  decinormal  permanganate  and  varying  proportions  of 
acid  during  five  days'  standing,  and  the  degree  of  decomposition 
of  the  permanganate  was  finally  determined  by  adding  a  small 
excess  of  oxalic  acid  to  the  mixtures  contained  in  Erlenmeyer 
flasks,  heating  to  about  80°  C.,  and  titrating  with  permanganate 
the  residual  oxalic  acid. 


Percentage 
H2SO?[i:i]. 

Time  elapsed. 

Percentage 
of  KMnO, 
decomposed. 

i  day. 

2  days. 

3  days. 

4  days. 

5  days. 

Color 

(     Color 

unchanged. 

1 

Color 

Color 

Color 

\  unchanged. 

Slight 

*»         1 

unchanged. 

unchanged. 

unchanged. 

Slight 

sediment. 

j.        3.0 

(^sediment. 

Slight 

scum. 

*>        { 

Color 
unchanged. 

Color 
unchanged. 

Color 
unchanged. 

!     Color 
unchanged. 
Slight 

Color 
unchanged. 
Slight 

7.4 

sediment. 

sediment. 

30         { 

Color 
unchanged. 

Color 
unchanged. 

Color 
unchanged. 

Reddish 
tinge. 
Slight 

Reddish 
tinge. 

-        6.9 

sediment. 

4o         | 

Color 
unchanged. 

Color 
unchanged. 

f    Tinged 
•(    with  red- 
l^dish  brown 

1  Reddish 
j      brown. 

Reddish 
brown. 

-       39-2 

( 

Color 

Color 

Reddish 

Reddish 

Red 

50         j 

unchanged. 

unchanged. 

i     brown. 

brown. 

brown. 

^        57-4 

, 

Color 

Color 

Reddish 

Sherry 

Reddish 

) 

€o         j 

redder. 

redder. 

brown. 

brown. 

olive. 

|       58.9 

( 

Color 

Color 

Sherry 

Reddish 

Reddish 

AY     -r 

70         ( 

redder. 

.  redder. 

brown. 

olive. 

olive. 

01.  1 

In  the  first  five  experiments  little  change  of  tint  was  noted 
upon  the  addition  of  the  oxalic  acid  to  the  cold  solution,  but  in 
the  last  two  experiments  the  reddish-olive  color  became  at  once 
distinctly  red  —  presumably  because  the  higher  sulphate  of  man- 
ganese was  attacked  in  the  cold  by  the  oxalic  acid  (as  Brauner 
has  shown) ,  and  so  the  natural  color  of  the  permanganate  was 
permitted  to  assert  itself.  The  extreme  decomposition  —  that 
which  took  place  in  the  last  experiment,  in  which  70  per  cent  of 
the  [i  :  i]  acid  was  present  —  corresponds  nearly  to  the  reduc- 
tion of  the  entire  amount  of  permanganate  present  to  the 
condition  of  oxidation  of  MnO2  which  is  known  to  exist  in  com- 
bination with  sulphuric  acid  in  the  form  of  a  higher  manganic 
sulphate.  It  is  to  be  noted  that  the  separation  of  the  insoluble 
higher  oxide  took  place  only  when  the  percentage  of  acid  was  low. 


APPLIANCES   AND    GENERAL   PROCEDURE 


45 


In  still  another  series  of  experiments  the  solution  of  perman- 
ganate was  mixed  with  sulphuric  acid  previously  diluted  with  an 
equal  volume  of  water,  and  cooled;  and  after  the  lapse  of  time 
indicated  oxalic  acid  was  added  in  quantity  a  little  more  than 
sufficient  to  bleach  the  permanganate.  The  solution  was  warmed 
to  about  80°  C.,  and  the  residual  oxalic  acid  titrated  by  gradual 
addition  of  more  permanganate.  The  difference  between  the 
amount  of  permanganate  needed  under  the  conditions  to  destroy 
the  known  amount  of  oxalic  acid,  and  that  used  in  the  deter- 
mination of  the  standard,  should  measure  the  oxygen  lost  and  the 
permanganate  decomposed  under  the  action  of  the  sulphuric  acid. 
The  results  and  details  of  these  experiments  are  given  below. 


Residual  Permanganate  Reduced  by  Oxalic  Acid  at  80°. 


H2S04  [i  :  ij. 
cm.3 

Water, 
cm.  3 

KMnO4  in  deci- 
normal  solution. 

cm.3 

Percentage  of 
H,S04li:i). 
in  solution 
during  action. 

Percentage  of 
KMnO4 
decomposed.. 

A.    Treated  immediately. 


2' 

8 

IO 

IO 

o 

4 

6 

IO 

20 

o 

6 

4 

IO 

30 

•  °  5 

8 

2 

10 

40 

1.6 

10 

IO 

50 

i.  9 

B.    Treated  after  standing  eight  hours  at  ordinary  temperature. 


2 

8 

IO 

10 

-0.3 

4 

6 

IO 

20 

—0.3 

6 

4 

10 

30 

!-3 

8 

2 

10 

40 

5-3 

IO 

10 

50 

15  7 

C.   Treated  after  standing  five  days  at  ordinary  temperature. 


2 

8 

10 

IO 

4  o 

4 

6 

10 

20 

21  6 

6 

4 

1C 

30 

49  7 

8 

2 

IO 

40 

55  9 

IO 

IO 

50 

564 

D.   Treated  after  standing  one  and  one-half  hours  at  8o°-9o°  C. 


2 

8 

10 

IO 

i  3 

4 

6 

IO 

20 

43-8 

6 

4 

IO 

3° 

35  9 

8 

2 

IO 

40 

49.1 

10 

10 

50 

55-3 

46 


METHODS  IN  CHEMICAL  ANALYSIS 


It  is  obvious  that  the  decomposition  of  the  permanganate  in- 
creases directly  in  each  series  of  experiments  with  the  increase 
in  the  proportion  of  sulphuric  acid,  that  the  amount  of  decom- 
position is  greater  as  the  time  of  action  is  extended,  and  that 
increase  of  temperature  heightens  the  change.  It  is  noted  in 
particular  that  the  presence  of  ten  per  cent  of  [i  :  i]  sulphuric 
acid  induces  at  the  ordinary  temperature  no  immediate  decom- 
position of  the  permanganate,  none  in  eight  hours,  and  a  break- 
ing down  amounting  to  four  per  cent  in  five  days;  and  that  the 
presence  of  fifty  per  cent  of  acid  of  the  same  strength  occasions 
the  decomposition  of  about  two  per  cent  at  once,  fifteen  per  cent 
in  eight  hours,  and  more  than  half  the  entire  amount  of  perman- 
ganate in  the  course  of  five  days.  It  is  evident,  also,  that  twenty 
per  cent  of  the  [i  :  i]  acid  produces  no  appreciable  effect  at 
ordinary  temperatures  and  under  exposures  of  a  few  hours  only. 
The  effect  of  heating  the  mixture  of  acid  and  permanganate  to 
So0  C.  for  an  hour  and  a  half  is  closely  comparable  with  that 
brought  about  by  the  five  days'  action  at  the  ordinary  tempera- 
ture. It  is,  of  course,  probable  that  some  decomposition  of  the 
permanganate  by  the  sulphuric  acid  is  brought  about  after  the 
addition  of  the  oxalic  acid  during  the  warming  of  the  mixture  up 
to  the  temperature  at  which  the  oxalic  acid  and  permanganate 

interact. 

Residual  Permanganate  Reduced  by  Ferrous  Sulphate. 


H2S04  [i  :  I), 
cm.* 

Water. 
cm.» 

KMnO4  in  deci- 
normal  solution. 

cm.* 

Percentage  of 
H2SO<|i:il 
in  solution 
during  action. 

Percentage  of 
KMn04 
decomposed. 

A.   Treated  at  once:  Volume,  20  cm.3 


2 

8 

10 

IO 

o.o 

4 

6 

IO 

20 

O.2 

6 

4 

IO 

3° 

O.I 

8 

2 

IO 

40 

O.  I 

10 

10 

50 

0-3 

B.   Treated  at  once:  Volume,  100  cm.3 


IO 

80 

10 

10 

O.  I 

20 

70 

10 

20 

O.I 

30 

60 

IO 

30 

o.o 

40 

SO 

IO 

40 

0-5 

50 
60 

40 
30 

10 
IO 

£ 

i-3 
3-o 

70 

20 

IO 

70 

5-o 

80 

IO 

IO 

80 

3-3 

90 

10 

90 

8.1 

APPLIANCES  AND   GENERAL   PROCEDURE  47 

Experiments  in  which  ferrous  sulphate  is  used  at  the  ordinary 
temperature  to  effect  the  reduction  of  the  residual  permanganate, 
instead  of  oxalic  acid  at  the  higher  temperature,  show  a  lower  de- 
gree of  decomposition,  as  is  natural.  The  increase  in  the  amount 
of  decomposition  as  the  proportions  of  sulphuric  acid  [i  :  i]  are 
advanced  beyond  50  per  cent  by  volume  is  striking.  The  results 
of  these  experiments  appear  in  the  accompanying  table. 

It  appears,  therefore,  that  when  potassium  permanganate  and 
sulphuric  acid  are  brought  into  solution  together  under  these  con- 
ditions, there  is  likelihood  of  a  reduction  of  the  former,  which  is 
greater  as  the  strength  of  the  acid  is  increased,  as  the  tempera- 
ture is  raised,  and  as  the  duration  of  action  is  extended.  It  ap- 
pears further,  at  least  when  the  acid  is  not  present  in  proportion 
greater  than  50  per  cent  of  the  [i  :  i]  mixture,  that  in  the  early 
stages  of  the  action  the  oxygen  lost  to  the  permanganate  is  liber- 
ated, and  that  later  on  the  decomposition  of  the  permanganate 
results  in  the  precipitation  of  manganese  as  a  higher  oxide  or  in 
the  formation  of  a  higher  sulphate.  The  first  effect  of  the  mutual 
action  of  the  acid  and  the  permanganate  is  to  set  free  perman- 
ganic acid,  which,  being  unstable,  breaks  up  with  the  results 
described. 

The  bearing  of  these  observations  and  inferences  upon  the 
question  of  the  action  of  potassium  permanganate  during  oxida- 
tions carried  on  in  the  presence  of  sulphuric  acid  is  obvious ;  for, 
if  the  aqueous  acid  is  able  to  liberate  permanganic  acid  in  such 
proportions  as  to  be  spontaneously  unstable,  it  is  reasonable  to 
presume  that  any  reducing  substance  present  at  the  time  of  such 
action  may,  by  virtue  of  its  attractive  action  upon  the  oxygen 
of  many  more  molecules  of  the  permanganic  acid  than  would 
be  necessary  to  supply  the  exact  amount  needed  for  perfect  oxi- 
dation, tend  to  increase  the  general  instability  of  the  already 
unstable  molecules  and  so  set  up  a  far-reaching  decomposition. 
These  considerations  throw  light  upon  the  phenomena  observed 
by  Brauner*  in  the  oxidation  of  tellurous  oxide  in  presence  of 
sulphuric  acid;  and  the  fact  that  the  liberation  of  free  oxygen 
in  this  special  ca.se  is  more  noticeable  than  in  the  oxidation  of 
ferrous  salts  or  oxalic  acid,  for  example,  is  explicable  in  the  light 
of  Brauner's  observation  that  the  attraction  of  tellurous  oxide 
for  oxygen  is  greatly  inferior  to  that  of  these  substances  for 
*  Jour.  Chem.  Soc.,  1891,  238. 


48  METHODS  IN  CHEMICAL  ANALYSIS 

oxygen  —  not  sufficient,  in  fact,  to  break  up  so  unstable  a  sub- 
stance as  manganic  sulphate,  which  is  at  once  reduced  by  ferrous 
salts  or  oxalic  acid. 

The  practical  lesson  to  be  drawn  from  the  investigation  is  the 
desirability  of  keeping  the  acid  present  in  oxidations  effected  by 
the  agency  of  permanganate  at  the  lowest  limit  consistent  with 
perfect  oxidation;  the  time  of  digestion  with  an  excess  of  per- 
manganate as  small  as  may  be;  and  the  temperature,  if  possible, 
not  above  the  ordinary  room  temperature. 

As  to  the  correlative  question  of  the  liberation  of  oxygen  dur- 
ing oxidations  by  potassium  permanganate  in  alkaline  solution, 
experience  in  the  collection  of  the  gas  liberated  in  oxidations 
effected  in  presence  of  acid  leads  to  distrust  of  the  evidence  of 
such  experiments  unless  the  amount  of  gas  liberated  is  consider- 
able. While,  on  the  one  hand,  small  quantities  of  liberated  gas 
may  be  so  completely  absorbed  as  not  to  appear  free  at  all,  it 
often  happens,  on  the  other  hand,  that  the  simple  admixture  of 
unlike  liquids — such,  for  example,  as  a  solution  of  potassium  per- 
manganate with  sulphuric  acid  of  strength  insufficient  to  liberate 
oxygen  —  may  bring  about  a  very  appreciable  liberation  of  dis- 
solved gases.  So  far  as  appears,  however,  the  affirmation  that 
oxygen  is  liberated  in  oxidations  by  potassium  permanganate  in 
alkaline  solutions  rests  upon  evidence  of  that  nature  only. 
Hydrochloric  Lowenthal  and  Lenssen*  were  the  first  to  show 
Acid  with  Fer-  that  the  titration  of  a  ferrous  salt  by  potassium  per- 
manganate in  the  presence  of  hydrochloric  acid, 
according  to  the  process  of  Margueritte,t  is  vitiated  by  the  evolu- 
tion of  chlorine  outside  the  main  reaction,  and  to  point  out  that 
a  remedy  for  the  difficulty  is  to  be  found  in  the  titration  of  the 
ferrous  salt  in  divided  portions,  other  equal  volumes  of  the  ferrous 
solution  being  added  to  the  liquid  in  which  the  first  titration  is 
accomplished  until  the  amount  of  iron  indicated  by  successive 
titrations  becomes  constant.  KesslerJ  showed  the  restraining 
influence  of  certain  sulphates,  of  manganous  sulphate  in  par- 
ticular, upon  the  irregular  and  undesirable  interaction  of  the 
permanganate  and  hydrochloric  acid,  and  Zimmermann,§  in 

*  Zeit.  anal.  Chem.,  i,  329. 

t  Ann.  Chim.  Phys.,  [3],  xviii,  244. 

f  Ann.  Phys.  Chem.,  cxviii,  48;   cxix,  225-226. 

§  Ann.  Chem.,  ccxiii,  302. 


APPLIANCES  AND   GENERAL  PROCEDURE  49 

apparent  ignorance  of  Kessler's  forgotten  proposal,  advocated  the 
introduction  of  a  manganous  salt,  best  the  sulphate,  into  the 
ferrous  salt  to  be  determined,  thus  accomplishing  the  purpose  of 
the  empirical  procedure  of  Lowenthal  and  Lenssen.  The  pro- 
tective influence  of  the  manganous  salt  turns  apparently,  as 
Zimmermann  suggested,  upon  the  initiation  of  Guyard 's  reaction, 
according  to  which  the  permanganate  and  manganous  salt  in- 
teract to  form  a  higher  oxide  of  manganese  capable  of  oxidizing 
the  ferrous  salt,  but  slow  to  act  upon  the  hydrochloric  acid.* 
According  to  Volhard,f  the  reaction  of  Guyard  is  favored  and 
hastened  by  heat  and  concentration  of  the  solution,  while  it  is 
delayed  by  acidity  and  dilution ;  but  even  in  solutions  containing 
very  little  manganous  salt  and  a  considerable  quantity  of  free 
acid,  the  faint  rose  color  developed  by  the  careful  addition  of 
permanganate  ultimately  vanishes  until  every  trace  of  the  man- 
ganous salt  is  precipitated.  When  a  considerable  amount  of  the 
salt  is  present  interaction  follows  immediately  the  introduction 
of  the  permanganate. 

In  titrations  of  a  ferrous  salt  by  permanganate,  Zimmermann 
advocates  the  use  of  4  grm.  of  manganous  sulphate  uniformly. 
In  putting  this  matter  to  the  test,  Gooch  and  Peters {  have  found 
that  as  much  as  5  grm.  of  manganous  sulphate  may  be  present 
in  135  cm.3  of  the  liquid,  containing  about  5  cm.3  of  hydrochloric 
acid  of  full  strength,  without  interfering  with  the  regularity  of 
the  titration;  and  the  effect  is  trivial  even  when  the  amount  of 
manganous  sulphate  reaches  10  grm.  In  all  cases,  however, 
in  which  the  larger  amounts  of  manganous  salt  are  present,  the 
end  reaction  is  marked  by  the  advent  of  a  brownish-red  precipi- 
tate rather  than  the  clear  pink  of  the  soluble  permanganate ;  and 
it  is  obvious  that  in  case  a  substance  to  be  oxidized  were  not 
active  enough  to  act  with  rapidity  upon  the  product  of  the 
Guyard  reaction,  difficulty  might  follow  the  failure  to  adjust  the 
conditions  more  particularly. 

Regularity  of  action  is  also  noted  when  manganous  chloride  is 
substituted  for  the  sulphate,  and  in  this  respect  the  results  accord 
with  those  of  Zimmermann  and  differ  from  those  of  Wagner.  § 

*  See  also  Manchot,  Ann.  Chem.  (1902),  cccxxv,  105. 

t  Ann.  Chem.  (1879),  cxcviii,  318. 

t  F.  A.  Gooch  and  C.  A.  Peter?,  Am.  Jour.  Sci.,  [4],  vii,  461. 

§  Zeit.  physikal.  Chem.,  28,  33. 


METHODS  IN  CHEMICAL  ANALYSIS 


Total  volume 
at  beginning 
of  titration. 

HC1. 
(Sp.  gr.  1.09). 

FeCl2. 

KMnO4  H/IO. 

MnSO4.sH2O. 

MnCl2.4H2O. 

cm.* 

cm.8 

cm.1 

cm.J 

grm. 

grrn. 

135 

IO 

25 

21  .  70 

I 

135 

IO 

25 

21  .70 

3 

135 

10 

25 

21.70 

5 

135 

IO 

25 

21-75 

7 

135 

IO 

25 

21-75 

IO 

US 

2O 

25 

21-75 

10 

175 

5° 

25 

21-75 

IO 

135 

IO 

25 

11  .  70 

I 

135 

IO 

25 

21  .  70 

2 

145 

20 

25 

21  .70 

2 

155 

30 

25 

21-75 

3 

165 

40 

25 

2I.7O 

4 

Hydrochloric 
Acid  with 
Oxalic  Acid. 


It  has  been  stated*  that  hydrochloric  acid  inter- 
feres in  no  way  with  the  titration  of  oxalic  acid  by 
permanganate.  Gooch  and  Peters  find,  however, 
that  in  such  titrations  there  is  a  small  though  real  waste  of  per- 
manganate proportionate  to  the  amount  of  hydrochloric  acid 
present.  This  fact  is  brought  out  clearly  in  the  comparison  of 
experimental  results  in  the  following  table. 

Temperature  at  Beginning  about  80°  C. 


Approximate 
volume  at 
beginning  of 
titration. 

H2SO<  [i  :  i]. 

HC1. 
(Sp.  gr.  1.09.) 

Ammonium 
oxalate  n/io. 

KMn04. 

M/IO. 

Variation  from 
mean  of  A 
taken  as 
standard. 

cm.1 

cm.1 

cm.8 

cm.8 

cm.s 

cm.* 

A. 


200 

5 

50 

47-50 

o.oo 

2OO 

5 

.... 

50 

47-50 

o.oo 

200 

10 

.... 

50 

47-50 

o.oo 

2OO 

10 

.... 

5° 

47-5° 

0.00 

200 

25 

50 

47-50 

o.oo 

200 

25 



5° 

47-50 

o.oo 

B. 


150 

IO 

2.5 

25 

23.80 

+0.05 

15° 

IO 

2.5 

25 

23.90 

+0.15    ; 

150 

10 

5-o 

25 

23.90 

+0.15    I 

15° 

IO 

IO.O 

25 

24.00 

+0.25 

500 

5 

.... 

25 

23.80 

+0.05 

500 

IO 

IO.O 

25 

24.00 

+0.25 

500 

10 

IO.O 

25 

24.10 

+0.35 

*  Fleischer,  Volumetric  Analysis,  Trans.  Muir.,  p.  71 .   Zimmermann,  loc.  cit. 


APPLIANCES  AND    GENERAL   PROCEDURE 


Temperature  2o°-26°  C. 


Volume 
at  begin- 
ning of 
titration. 

H2SO4 
[i  :  i]. 

cm.* 

HC1. 
(Sp.  gr.  1.09) 

cm.* 

Ammo- 
nium 
oxalate, 

M/IO. 

cm.* 

KMnO4. 

W/IO. 

cm.* 

MnSO4.- 
5H20. 

grm. 

MnCl,.- 
4H20, 

grm. 

Variation 
from 
standard. 

cm.* 

IO 

23  oo 

OO4O 

-f-O.  15 

I  tO 

IO 

2C 

23  QO 

OI2O 

-f-O.  15 

IO 

2Z 

23.80 

+O.O5 

130 

IO 

2Z 

23  .  75 

.O4OO 

+O.OO 

130 

IO 

23  •  76 

.O5OO 

+O.OI 

130 

IO 

25 

23.  70 

.  IOOO 

—  0.05 

IO 

23    7$ 

2OOO 

o  oo 

I3O 

IO 

2? 

24   20 

O2OO 

-f-o  45 

I3O 

IO 

25 

23  95 

.O2OO 

-|-o.  20 

I3O 

IO 

23.80 

.0400 

I3O 

20 

25 

23  .  75 

.0400 

o.oo 

130 

3° 

25 

23.75 

.0400 

o.oo 

I3O 

IO 

25 

23.75 

I  .OOOO 

o.oo 

IO 

23    71? 

2    OOOO 

o  oo 

130 

IO 

2C 

23    7^ 

3  oooo 

o  oo 

I3O 

I 

23    72 

I    OOOO 

—  o  03 

I3O 

I 

2C 

23  .  74 

2  .OOOO 

—  O.OI 

I30 
130 

130 

I 
2 

3 

•• 

•>  10  10  »o 

d  <N  M 

23.72 
23.70 
23-75 

3.0000 
0.5000 
0.5000 

—0.03 
—0.05 
o.oo 

Temperature  about  80°. 


145 

IO 

IO 

2i? 

23  oo 

o  5000 

4-o  I  C 

145 

IO 

IO 

2< 

23  .  7O 

I  .OOOO 

—  o  OT 

500 

IO 

IO 

25 

23.7? 

I  .OOOO 

o.oo 

500 

IO 

25 

23.  70 

I  .OOOO 

—  0.05 

500 

IO 

25 

24.  10 

o  .  5000 

4-o.  35 

The  error  introduced  by  the  presence  of  hydrochloric  acid 
during  the  action  of  the  permanganate  upon  oxalic  acid  may  be 
obviated  by  the  introduction  of  a  manganous  salt,  but  the  per- 
sistence of  the  Guyard  reaction  is  liable  to  interfere  with  the  end 
reaction  of  oxidation  of  oxalic  acid  unless  an  adjustment  is. made 
between  the  quantity  of  the  manganous  salt,  the  amount  of  acid 
and  the  dilution.  It  appears,  also,  that  for  a  given  dilution  and 
strength  of  acid  less  manganous  salt  is  needed  in  a  cold  solution 
than  in  a  hot  solution.  Thus,  in  the  hot  solution,  at  a  dilution 
of  145  cm.3  to  500  cm.3  with  5  cm.3  of  strong  hydrochloric  acid, 
with  or  without  sulphuric  acid,  I  grm.  of  manganous  sulphate 
must  be  present;  while  in  the  cold  solution,  0.04  grm.  of  either 
the  sulphate  or  chloride  is  enough  to  secure  adequate  protective 


52  METHODS  IN   CHEMICAL   ANALYSIS 

effect.  Experience  showed,  however,  that  i.o  grm.  of  the  man- 
ganous  salt  should  be  present  in  order  to  push  the  reaction  with 
reasonable  speed,  and  that  this  amount  is  enough  to  so  affect  the 
conditions  of  equilibrium  that  titrations  in  moderate  volumes 
(100  cm.3  to  500  cm.3)  and  in  presence  of  hydrochloric  acid  (5  cm.3 
to  15  cm.3  of  the  strong  acid)  may  be  conducted  with  safety  and 
reasonable  rapidity,  either  with  or  without  sulphuric  acid,  at  the 
ordinary  atmospheric  temperature. 

Experimental  results  are  given  in  the  preceding  table. 
Effect  of  other  It  has  been  shown  in  the  preceding  account  that 
Chlorides.  when  potassium  permanganate  and  hydrochloric  acid 
react  at  high  concentrations  chlorine  is  evolved,  and  when  oxi- 
dations are  brought  about  by  permanganate  in  presence  of 
sulphuric  acid  and  small  amounts  of  chlorides,  the  tendency 
toward  evolution  of  chlorine  is  in  evidence.  This  fact  must  be 
taken  into  consideration  in  analytical  operations  involving  such 
conditions.  It  has  also  been  stated*  that  certain  chlorides  act 
catalytically  to  induce  further  decomposition  of  the  perman- 
ganate than  would  take  place  under  the  conditions  were  hydro- 
chloric acid  the  only  chloride  present;  but  Brown  f  has  shown 
that  this  assertion  is  a  mistake  due  to  faulty  analytical  procedure. 
When  a  definite  quantity  of  potassium  permanganate  is  digested 
with  a  definite  quantity  of  normal  hydrochloric  acid  under  de- 
nned conditions  of  time  and  temperature,  and  the  resulting  mix- 
ture is  treated  with  a  definite  amount  of  oxalic  acid,  the  excess  of 
which  is  determined  by  titration  with  permanganate,  the  data 
are  at  hand  for  calculating  the  amount  of  permanganate  appar- 
ently reduced  during  the  .digestion. 

Wagner  has  found  that  when  a  small  amount  of  n/io  ferric 
chloride  is  added  to  one  such  mixture  before  the  digestion  and  an 
equivalent  amount  of  n/io  hydrochloric  acid  to  another  such 
mixture,  more  permanganate  is  required  in  the  final  titration  of 
the  excess  of  oxalic  acid  in  the  mixture  which  contains  the  ferric 
salt  than  in  that  which  contains  no  ferric  salt,  a  fact  which  seems 
at  the  outset  to  imply  more  decomposition  of  permanganate 
when  the  digestion  takes  place  in  presence  of  the  ferric  salt. 
Brown  shows,  however,  that  when  the  chlorine  produced  in  the 
reaction  is  removed  by  a  current  of  carbon  dioxide  or  air  before 

*  Wagner,  Zeit.  physikal.  Chem.,  28,  33. 
f  James  Brown,  Am.  Jour.  Sci.,  [4],  xix,  31. 


APPLIANCES  AND    GENERAL  PROCEDURE 


53 


the  addition  of  the  oxalic  acid,  the  same  quantity  of  permanga- 
nate is  required  to  bring  about  the  final  coloration  whether  ferric 
chloride  is  present  or  not. 

The  conclusion  must  be  drawn,  then,  that  Wagner's  experi- 
ments in  no  way  show  the  catalytic  effect  of  ferric  chloride  in 
the  interaction  between  hydrochloric  acid  and  potassium  per- 
manganate, nor  do  they  furnish  evidence  in  support  of  the  as- 
sumed formation  of  chlor- ferrous  acid.  They  afford  simply  an 
indication  of  the  greater  or  less  retention  of  chlorine  in  solution, 
and  the  greater  or  less  oxidizing  action  of  this  chlorine  on  the 
oxalic  acid  in  the  presence  or  absence  of  ferric  chloride. 

The  following  tables  give  experimental  results. 

Digestion  without  the  Removal  of  Chlorine. 
(9.91  cm.3  H2C2O4=  20.25  cm-3  KMnO4.) 


n/i  HC1. 
cm.8 

w/ioHCl. 
cm.3 

M/IO 

FeCl3. 
cm.s 

KMnO4 
before 
digestion. 

cm.3 

Temper- 
ature. 

C°. 

Time  of 
digestion. 

min. 

«/IO 

H2C204. 
cm.3 

KMnO4 
to  color. 

cm.3 

KMnO 
appar- 
ently 
reduced 
during 
digestion. 

cm.3 

100 

9.91 

9.91 

5° 

60 

9.91 

15.89 

5-55 

IOO 

9.91 

9.91 

5° 

60 

9.91 

15-11 

4-77 

IOO 

9.91 

9.91 

50 

60 

9.91 

I5-I5 

4.81 

IOO 

9.91 

9.91 

5° 

60 

9.91 

15.07 

4-73 

IOO 

9.91 

9.91 

SO 

60 

9.91 

I5-I3 

4-79 

IOO 

9.91 

9.91 

5° 

60 

9.91 

!5-07 

4-73 

IOO 

9.91 

9.91 

50 

60 

9.91 

15.08 

4-74 

IOO 

9.91 

9.91 

5° 

60 

9.91 

15.02 

4.68 

IOO 

9.91 

9.91 

50 

60 

9.91 

14-85 

4.5i 

IOO 

9.91 

9.91 

5° 

60 

9.91 

14.40 

4.06 

IOO 

9.91 

9.91 

SO 

60 

9.91 

15-05 

4-7i 

IOO 

.... 

9.91 

9.91 

5° 

60 

9.91 

15.60 

5.26 

IOO 

9.91 

9.91 

50 

60 

9.91 

15-35 

5-oi 

IOO 

9.91 

9.91 

5° 

60 

9.91 

15-32 

4-98 

IOO 

9.91 

9.91 

50 

60 

9.91 

15.88 

5-54 

IOO 

9.91 

9.91 

50 

60 

9.91 

15.42 

5-o8 

IOO 

.... 

9.91 

9.91 

50 

60 

9.91 

15-45 

5-n 

IOO 

9.91 

9.91 

50 

60 

9.91 

15-95 

5-6i 

IOO 

9.91 

9.91 

50 

60 

9.91 

i5-4i 

5-07 

IOO 

9.91 

9.91 

50 

60 

9.91 

15-95 

5-6i 

IOO 

9.91 

9.91 

5° 

60 

9.91 

16.65 

6.31 

IOO 

.... 

9.91 

9.91 

5° 

60 

9.91 

15-75 

5-41 

IOO 

9.91 

9.91 

5° 

60 

9.91 

15-79 

5-45 

In  all  these  experiments  the  potassium  permanganate  was  en- 
tirely destroyed,  the  hydrochloric  acid  present  being  capable  of 
breaking  down  many  times  as  much  permanganate  under  iden- 
tical conditions  of  time  and  temperature. 


54 


METHODS  IN  CHEMICAL  ANALYSIS 


Chlorine  Removed  by  a  Current  of  Air  during  Digestion. 
(9.90  cm.3  approximately  w/io  H2C2C>4  =  20.09  cm.3  KMnO4.) 


J 

«4 

, 

1  .>>*> 

e 

t? 

3! 

:s 

fc 
75 

§'g 

o 
a 

10  HCI. 

10  FeCl3. 

ij 

£.! 

§ 

me  of  dige 

M| 

"d* 

-§0« 

•l"o  2 
o>  o  -H 

1 
II 

o; 

8 
o 

a 

35 

||| 

K 

^ 

la 

M 

o 

£ 

o 

a 

W 

M  ^ 

cm.3 

cm.* 

cm.8 

cm.3 

c°. 

min. 

cm.3 

cm.3 

cm.3 

IOO 

9.90 

9.90 

50 

60 

none 

none 

9.90 

18.88 

8.69 

IOO 

9.90 

9.90 

5° 

60 

none 

none 

9.90 

18.87 

8.68 

IOO 

9.90 

9.90 

50 

60 

none 

none 

9.90 

18.80 

8.61 

IOO 

9.90 

9.90 

50 

60 

none 

none 

9.90 

18.81 

8.62 

IOO 

9.90 

.... 

9.90 

50 

60 

none 

none 

9.90 

18.80 

8.61 

IOO 

9.90 

9.90 

50 

60 

none 

none 

9.90 

18.82 

8.63 

IOO 

9.90 

9.90 

5° 

30 

none 

none 

9.90 

18.77 

8.58 

IOO 

9.90 

9.90 

50 

3° 

none 

doubtful 

9.90 

18.70 

8-51 

IOO 

9.90 

9.90 

50 

15 

none 

very  faint 

9.90 

18.70 

8-51 

IOO 

9.90 

9.90 

50 

15 

none 

very  faint 

9.90 

18.68 

8-49 

IOO 

9.90 

9.90 

50 

60 

none 

none 

9.90 

18.87 

8.68 

IOO 

9.90 

9.90 

50 

60 

none 

none 

9.90 

18.85 

8.66 

IOO 

9.90 

9.90 

50 

60 

none 

none 

9.90 

18.81 

8.62 

IOO 

9.90 

9.90 

50 

60 

none 

none 

9.90 

18.81 

8.62 

IOO 

9.90 

9.90 

50 

60 

none 

none 

9.90 

18.87 

8.68 

IOO 

9.90 

9.90 

50 

60 

none 

none 

9.90 

18.85 

8.66 

IOO 

9.90 

9.90 

5° 

30 

none 

none 

9.90 

18.80 

8.61 

IOO 

9.90 

9.90 

3° 

none 

doubtful 

9.90 

18.72 

8-53 

IOO 

9.90 

9.90 

50 

15 

none 

very  faint 

9.90 

18.65 

8.46 

IOO 

9.90 

9.90 

50 

15 

none 

very  faint 

9.90 

18.67 

8.48 

With  cadmium  chloride  and  gold  chloride,  the  apparently 
catalytic  effect  is  due  entirely  to  chlorine  retained  in  solution, 
while  with  chromic  chloride  and  platinic  chloride  increased  effects 
are  due  partly  to  chlorine  retained  in  solution  and  partly  to  the 
total  reduction  of  residual  oxides  of  manganese. 


ACIDIMETRY   AND   ALKALIMETRY. 

The  Use  of  Succinic  Acid  as  the  Standard  in  Neutralization 

Processes. 

Methods  for  preparation  and  the  use  of  succinic  acid  as  a  stand- 
ard in  the  titration  of  an  alkali  hydroxide  have  been  studied  by 
Phelps  and  Hubbard.*  Succinic  acid  was  prepared  by  hydrolysis 
of  the  pure  ethyl  ester,  by  hydration  of  the  pure  anhydride,  by 
recrystallization  of  the  commercial  acid  from  hot  water,  and 
by  recrystallization  of  the  commercial  acid  from  hot  water 
*  I.  K.  Phelps  and  J.  L.  Hubbard,  Am.  Jou'-.  Sci.,  [4],  xxiii,  211. 


APPLIANCES  AND   GENERAL  PROCEDURE 


55 


containing  nitric  acid.  The  product  obtained  by  each  of  these 
methods  dries  to  a  constant  weight  in  air,  and  loses  nothing  on 
further  standing  over  sulphuric  acid. 

From  ethyl  succinic  ester,  boiling  at  2I3.3°-2I3.5°  C.  under  a 
barometric  pressure  of  749  mm.,  pure  succinic  acid  was  obtained 
by  boiling  it  for  four  hours  on  a  return  condenser  with  nitric 
acid  and  water  in  these  proportions:  20  cm.3  of  succinic  ester, 
200  cm.3  of  water,  three  drops  of  nitric  acid.  The  solution  was 
evaporated  to  crystallization,  and,  after  filtering  off  from  the 
mother  liquor,  the  solid  product  was  recrystallized  from  distilled 
water.  The  crystals  dried  carefully  in  the  open  air  to  constant 
weight  melted  in  an  open  capillary  tube  at  182.3°. 

Ammonium  Hydroxide  against  Succinic  Acid. 


Succinic  acid, 
grm. 

HC1  value  of  succinic  acid. 

Found, 
grm. 

Theory, 
grm. 

Error, 
grm. 

Acid  from  ester. 


0.2000 

0.1235 

0.1235 

o.oooo 

O.  2OOO 

0.1235 

0.1235 

o.oooo 

O.2OOO 

O.I23S 

0.1235 

o.oooo 

O.2OOO 

0.1235 

0.1235 

0.0000 

Acid  from  anhydride. 


0:2000 

O.I23I 

0.1235 

—  O.OOO4 

0.2000 

0.1233 

0.1235 

—  O.OOO2 

O.2OOO 

0.1234 

0.1235 

—  O.OOOI 

O.2OOO 

0.1234 

0.1235 

—  o.oooi 

Acid  recrystallized  from  water. 


O.2OOO 

O.I23I 

0.1235 

—  O.OOO4 

O.2OOO 

O.I23I 

0-1235 

—  O.OOO4 

0.2000 

O.I23I 

0.1235 

—  0.0004 

0.2000   , 

O.I23I 

0.1235 

—  O.OOO4 

O.  2OOO 

O.I23I 

0.1235 

—  O.OOO4 

O.2OOO 

O.I23I 

0.1235 

—  O.OOO4 

Acid  recrystallized  from  water  containing  nitric  acid. 


O.2OOO 
0.2000 

0.1233 
0.1233 

0.1235 
0.1235 

—  O.OOO2 
—  O.OOO2 

For  the  preparation  of  the  acid  from  succinic  anhydride  the 
anhydride  was  purified  by  recrystallizing  from  absolute  alcohol 


56  METHODS  IN   CHEMICAL  ANALYSIS 

until  the  crystals  obtained,  after  carefully  drying,  melted  sharply 
at  H9°C.  The  pure  anhydride  obtained  in  this  manner  was 
converted  to  the  acid  by  dissolving  it  in  distilled  water  heated 
to  the  boiling  point.  The  crystals  formed  on  cooling  the  solution 
were  filtered  off  and  dried  in  air  and  over  sulphuric  acid.  The 
melting  point  of  the  product  was  182.8°. 

Preparations  made  by  dissolving  the  acid  of  commerce  in  dis- 
tilled water  at  the  boiling  point,  crystallizing  by  cooling,  and 
drying  in  the  air,  melted  at  181.7°.  The  acid  made  by  dissolv- 
ing the  succinic  acid  of  commerce  in  boiling  water,  adding  nitric 
acid,  crystallizing,  and  drying,  melted  at  182.3°. 

Tests  of  these  preparations  of  succinic  acid  were  made  by 
titrating  solutions  of  weighed  amounts  by  approximately  n/io 
ammonium  hydroxide,  with  cochineal  as  an  indicator,  the  ammo- 
nium hydroxide  having  been  previously  standardized  against  ap- 
proximately n/io  hydrochloric  acid,  the  exact  strength  of  which 
had  been  determined  gravimetrically  by  precipitation  with  special 
precautions  by  silver  nitrate.  The  results  are  given  on  page  55. 

Organic  Acids  and  Acid  Anhydrides  as  Standards  in  Neutraliza- 
tion Processes. 

Phelps  and  Weed*  have  shown  that,  with  phenolphthalein  as 
an  indicator,  pure  sodium  hydroxide  in  solution  and  pure  barium 
hydroxide  in  solution  may  be  determined  very  exactly  by  titra- 
tion  against  weighed  amounts  of  succinic  acid,  succinic  an- 
hydride, malonic  acid,  benzoic  acid,  phthalic  acid  and  phthalic 
anhydride  used  as  standards.  Following  are  experimental  results 

(PP-  57,  58). 

Phelps  and  Weed  point  out  that  these  organic  acids  and  acid 
anhydrides,  in  pure  state,  make  standards  in  alkalimetry  and 
acidimetry,  as  accurate  as  the  best  previous  standard,  —  hydro- 
chloric acid  determined  gravimetrically  as  silver  chloride.  The 
most  serviceable  are  those  most  readily  soluble  in  water,  —  suc- 
cinic and  malonic  acids. 

Phelps  and  Weedf  point  out,  also,  that  these  organic  acids  and 
anhydrides  —  since  they  may  be  used  to  fix  the  standards  of 
alkali  hydroxides,  and  these  to  fix  the  standard  of  hydrochloric 

*  I.  K.  Phelps  and  L.  H.  Weed,  Am.  Jour.  Sci.,  [4],  xxvi,  138. 
t  Am.  Jour.  Sci.,  [4],  xxvi,  143. 


APPLIANCES  AND   GENERAL  PROCEDURE 


57 


Sodium  Hydroxide  and  Barium  Hydroxide  against  Succinic  Acid  and  Succinic 

Anhydride. 


Succinic  acid. 

grm. 

Succinic 
anhydride. 

grm. 

HC1  value  of 
NaOH  used. 

grm. 

HCl  value  of 
BaO2H2  used. 

grm. 

Theory  in 
terms  of  HCl. 

grm. 

Error  in 
terms  of  HCl. 

grm. 

fo    2OOO 

O   1236 

O.I23S 

+O.OOOI 

1  o    2OOO 

o  1238 

o.  1235 

+0.0003 

I  -\  O    2OOO 

O    1237 

o.  1235 

+O  .  OOO2 

O    2OOO 

o.  1236 

0.1235 

+O.OOOI 

|^O    2OOO 

o.  1236 

o.  1235 

-f-o.oooi 

fo    2OOO 

O    1237 

O.  123"? 

+O.OOO2 

2   •{  O    2OOO 

O    1237 

O.  I23i? 

+O.OOO2 

1  O    2OOO 

O    1237 

O.I235 

+O.OO02 

t  O    2OOO 
i  u  .  tvw 

o  1237 

O.I235 

+O.OOO2 

0    (  O    2OOO 

o.  1237 

O.I235 

+O  .  OOO2 

f  O  .  2OOO 
1  O  .  2OOO 
|  O.2OOO 
{o  .  2OOO 

O    2OOO 

O    14.^8 

0.1238 
0.1237 
0.1235 
0.1236 

0.1235 
0.1235 
0-1235 
0.1235 

o.  1458 

+o  .  0003 

+O  .  OOO2 

o.oooo 

+O.OOOI 

o.oooo 

o   2000 

o  1458 

o.  1458 

o.oooo 

O    2OOO 

o  14.^0 

o.  1458 

-J-O.OOOI 

O.  2OOO 

o.  1458 

0.1458 

o.oooo 

O.  2OOO 

o.  1457 

0.1458 

—  O.OOOI 

O.  2OOO 

o.  1456 

o.  1458 

—  O.OOO2 

O.  2OOO 

o.  1459 

0.1458 

+O.OOOI 

O    2OOO 

O    14^8 

o  1458 

o  oooo 

1.  Freshly  made  the  ester  and  dried  over  sulphuric  acid. 

2.  Dried  for  a  year  over  sulphuric  acid. 

3.  Dried  for  a  year  over  calcium  chloride. 


Sodium  Hydroxide  and  Barium  Hydroxide  against  Malonic  Acid* 


Malonic  acid, 
grm. 

HCl  value  of 
NaOH  used. 

grm. 

HCl  value  of 
BaO2H2  used. 

grm. 

Theory  in  terms 
of  HCl. 

grm. 

Error  in  terms 
of  HCl. 

grm. 

O.20OO 

o.  1404 

o.  1402 

-}-O.OOO2 

O.  2OOO 

O.  1403 

o.  1402 

-J-O.OOOI 

O.  2OOO 

o  1402 

o  1402 

o  oooo 

O.  2OOO 

o  1401 

o  1402 

—  o  oooi 

O.2OOO 

o  1401 

o  1402 

—  o  oooi 

O.2OOO 

o  1400 

o  1402 

—  O   OOO2 

O.2OOO 
O.  2OOO 



o.  1402 
o  1400 

0.1402 
o  1402 

o.oooo 

—  O   OOO2 

*  Prepared  from   malonic  ester  by  hydrolysis,  recrystallized  from  water,  and  dried  over 
sulphuric  acid. 


METHODS  IN  CHEMICAL  ANALYSIS 


Sodium  Hydroxide  and  Barium  Hydroxide  against  Benzoic  Acid.* 


Benzoic  acid, 
grm. 

HC1  value  of 
NaOH  used. 

grm. 

HC1  value  of 
BaO2H2  used. 

grm. 

Theory  in  terms 
of  HC1. 

grm. 

Error  in  terms 
of  HC1. 

grm. 

O.2OOO 

o  2000 

0.0598 
O.O5Q9 

0.0597 
0.0597 

+O.OOOI 
+O.OOO2 

O.  2OOO 
O.  2OOO 

0.0597 
0.0598 



0.0597 
0.0597 

0.0000 
+O.OOOI 

O    2OOO 

o  otjoS 

O  o<07 

+o  oooi 

O    2OOO 

o  o<Q7 

0.0507 

o.oooo 

O    2OOO 

0.0597 

0.0597 

o.oooo 

O.2OOO 



0.0597 

0.0597 

o.oooo 

*  Prepared  by  treating  benzoic  ester  with  sodium  hydroxide,  acidifying  with  hydrochloric  acid, 
and  twice  crystallizing  from  water  the  precipitate,  and  drying  over  sulphuric  acid.  In  these 
experiments,  alkali  in  amount  nearly  sufficient  for  neutralization  was  run  into  the  flask  upon  the 
acid,  which  was  then  heated  to  bring  about  solution  of  the  acid. 


Sodium  Hydroxide  and  Barium  Hydroxide  against  Phthalic  Acid*  and  Phthalic 

Anhydride.^ 


Phthalic  acid, 
grm. 

Phthalic 
anhydride 

grm. 

HC1  value  of 
NaOH  used. 

grm. 

HC1  value  of 
BaO2H2  used. 

grm. 

Theory  in 
terms  of  HC1. 

grm. 

Error  in  terms 
of  HC1. 

grm. 

O    2OOO    • 

O  0880 

0.0878 

+0.0002 

O    2OOol 

0  0880 

0.0878 

+0.0002 

O    2OOO 

0.0879 

0.0878 

+O.OOOI 

0.2000 
O    2OOO 



0.0878 

0  0876 

0.0878 
0.0878 

0  .  OOOO 
—  O.O002 

O.  2OOO 
O.2OOO 
0.2000$ 

O.  2OOO 

o  .  0986 

0.0877 
0.0878 
0.0879 

0.0878 
0.0878 
0.0878 
o  .  0985 

—  O.OOOI 
O  .  OOOO 
+O.OOOI 
+O.OOOI 

O    2OOO 

o  0085 

o  .  0985 

O  .  OOOO 

O.2OOO 

0.0986 

o  .  0985 

+0.0001 

O.  2OOO 

0.0987 

0.0985 

+O.OO02 

O.2OOO 

o  .  0986 

0.0985 

+O.OOOI 

O    2OOO 

o  .  0985 

0.0985 

O  .  OOOO 

O    2OOO 

0.0986 

0.0985 

+O.OOOI 

O    2OOO 

o  0987 

0.0985 

+  O.OOO2 

*  Prepared  from  the  commercial  anhydride  by  dissolving  in  boiling  water,  crystallizing  from 
the  filtered  solution  by  cooling,  and  drying  finally  over  sulphuric  acid, 
t  Prepared  by  distilling  in  vacua  the  phthalic  anhydride  of  commerce. 
J  Titrated  at  the  room  temperature ;  all  others  titrated  after  heating  to  secure  solubility. 


APPLIANCES  AND  GENERAL  PROCEDURE        59 

acid  or  sulphuric  acid,  which  in  turn  (under  definite  conditions 
of  concentration,  as  will  be  shown*)  may  set  free  from  an  iodide- 
iodate  mixture  an  exactly  equivalent  amount  of  iodine  —  may  be 
used  indirectly  to  set  the  standards  used  in  iodometric  analysis. 

The  Use  of  the  lodide-Iodate  Mixture  and  the  Estimation  of  the 

Iodine  Evolved. 

Determination  It  is  well  known  that  when  a  free  mineral  acid  is 
of  Free  Acids.  a(ided  to  a  neutral  mixture  of  metallic  iodate  and 
iodide,  the  iodate  is  reduced  and  iodine  is  liberated  according  to- 
the  equation: 

RI03  +  5  RI  +  3  H2S04  =  3  I2  +  3  R2SO4  +  3  H2O. 
This  reaction  is  complete  and  nonreversible  under  suitable  con- 
ditions and  may  therefore  be  applied  to  the  estimation  of  amounts 
of  iodate,  iodide  or  mineral  acid  present  in  an  unknown  solution. 
A  solution  of  iodate  to  be  analyzed  is  mixed  with  an  excess  of 
iodide  and  mineral  acid,  the  resulting  free  iodine  estimated  by 
directly  titrating  with  sodium  thiosulphate  or  arsenious  acid  after 
neutralization  and  one-sixth  of  the  amount  found  taken  as  equiv- 
alent to  the  iodate  originally  present. f  Similarly,  a  solution  of 
iodide  to  be  analyzed  is  mixed  with  an  excess  of  iodate  and  min- 
eral acid,  the  resulting  free  iodine  estimated  by  directly  titrat- 
ing in  alkaline  solution  with  arsenious  acid,  and  five-sixths  of  its 
amount  taken  as  equivalent  to  the  iodide  originally  present. £ 
The  solution  of  mineral  acid  to  be  analyzed  is  mixed  with  an  ex- 
cess of  iodate  and  iodide,  the  resulting  free  iodine  estimated  by 
directly  titrating  with  sodium  thiosulphate,  and  its  entire  amount 
taken  as  equivalent  to  the  amount  of  mineral  acid  originally 
present.§  Groger  has  applied  the  last- mentioned  method  to  the 
direct  analysis  of  various  mineral  acids  and  has  also  indirectly 
analyzed  solutions  of  alkali  hydroxides  and  carbonates  by  adding: 
the  solution  to  be  analyzed  to  a  measured  volume  of  mineral 
acid,  previously  standardized  by  the  above  method,  and  esti- 
mating the  small  excess  of  free  mineral  acid  that  finally  remains 

*  See  page  60. 

t  Rammelsberg,  Ann.  Phys.,  cxxxv,  493;  Walker,  Am.  Jour.  Sci.,  [4],  iv> 
235- 

J  Gooch  and  Walker,  cf.  page  454. 

§  Kjeldahl,  Zeit.  anal.  Chem.,  xxii,  366;  Furry,  Am.  Chem.  Jour.,  vi,  341; 
Groger,  Zeit.  angew.  Chem.,  1894,  52. 


6o 


METHODS  IN  CHEMICAL  ANALYSIS 


by  the  same  method.  The  only  difficulty  with  the  Groger 
process  lies  in  the  fact  that  in  dilute  solutions,  as  shown  by 
Furry,*  the  end-point  of  the  final  reaction  between  iodine  and 
sodium  thiosulphate  is  somewhat  obscured  by  a  peculiar  back- 
play  of  color  due  to  a  continuous  slow  liberation  of  iodine  in 
the  system. 

When  the  suitably  concentrated  iodide-iodate  mixture  (i  grm. 
of  KI  and  0.166  grm.  of  KIO3  in  50  cm.3)  is  acted  upon  by 
a  definite  quantity  of  approximately  decinormal  hydrochloric 
acid  in  a  total  volume  not  exceeding  100  cm.3,  the  iodine  set 
free,  as  determined  by  titration  with  n/io  sodium  thiosulphate 
(standardized  against  iodine  of  strength  fixed  by  titration  against 
weighed  arsenious  oxide),  measures  very  exactly,  according  to 
Phelps  and  Weed,f  the  acid  taken.  The  recorded  results  of 
experiments,  in  which  the  end  reaction  was  brought  about  by 
iodine  added  directly  or  set  free  by  addition  of  more  standard 
acid  after  the  bleaching  by  thiosulphate,  are  exceedingly  good. 

Volume  about  100  cm.3 
HC1  values. 


HC1  solution 
used. 

gnu. 

NajSjOa  solution 
used. 

grm. 

Iodine  solution 
used. 

grm. 

HC1  found, 
grm. 

Error  in  HC1. 
grm. 

o.  1074 

O    IO7< 

O   IO7< 

-J-O  .  OOOI 

o.  1074 

w  •  -*-w  /  O 

o.  1074 

\J  .    A  v/  ^ 

o.  1074 

O.  OOOO 

0.1074 

0.1075, 

0.1075 

+0.0001 

0.0520 

0.0520 

O.O52O 

o.oooo 

0.1560 

0.1562 



0.1562 

+O.OO02 

0.0645 

0.0712 

0.0068 

o  .  0644 

—  O.OOOI 

0.0968 

0.1068 

O.OIOI 

0.0967 

—  0.0001 

0.0645 

0.0712 

0.0067 

o  .  0645 

o.oooo 

0.0968 

0.1068 

O.OIOI 

0.0967 

—  O.OOOI 

0.0484 

0-0534 

o  .  0050 

o  .  0484 

o.oooo 

0.0645 

0.0748 

0.0104 

0.0644 

—  O.OOOI 

0.0484 

0.0534 

o  .  0050 

o  .  0484 

o.oooo 

0.0645 

0.0712 

0.0066 

o  .  0646 

+O.OOOI 

Determination  Alkali  hydroxides  may  be  determined  by  the 
Hyllroxides  and  acti°n  of  an  excess  of  standard  hydrochloric  acid  or 
Carbonates.  sulphuric  acid,  the  excess  being  determined  by  esti- 
mation of  the  iodine  set  free  by  the  iodide-iodate  mixture  at  suit- 
able concentrations.  Alkali  carbonates  may  also  be  similarly 

*  Am.  Chem.  Jour.,  vi,  341. 
t  Am.  Jour.  Sci.,  [4],  xxvi,  143. 


APPLIANCES   AND    GENERAL  PROCEDURE 


6l 


determined  by  treatment  with  an  excess  of  sulphuric  acid,  boil- 
ing to  remove  carbon  dioxide  from  the  solution  containing  the 
nonvolatile  acid,  and  determination  of  iodine  liberated  by  the 
action  of  the  iodide-iodate  mixture  in  a  total  volume  of  about 
100  cm.3  Results  obtained  by  this  treatment  of  sodium  hy- 
droxide first  treated  with  carbon  dioxide  are  given  in  the  follow- 
ing statement,  and,  for  comparison,  results  obtained  by  titration 
of  the  excess  of  standard  acid  by  standard  sodium  hydroxide. 

Volume  not  Exceeding  100  cm.3 


HC1  values. 

Treatment 
with  CO2. 

min. 

NaOH 
solution 
used. 

grm. 

H2S04 
solution 
used. 

gnu. 

NaOH 
solution  to 
coloration. 

grm. 

Na,S,0, 

solution 
used. 

grm. 

Iodine 
solution  to 
coloration. 

grm. 

Difference  in 
terms  of  HC1. 

grm. 

je 

O    IOI2 

o.  1306 

O.O2Q3 

-|-O  OOOI 

•3Q 

O    IOI2 

o.  1219 

0.0205 

-|-O  OOO2 

jr 

O    1349 

o.  1524 

0.0172 

-|-o  0003 

2Q 

O    I  34Q 

o  1568 

o  0216 

-f"O   OOO^ 

T  I? 

O    IOI2 

o  1306 

O  0302 

O  0013 

-f-o  0005 

•2Q 

O    IOI2 

o  1306 

o  0302 

O  0013 

-r~O   OOO^ 

15 

35 

0.1349 
0.1349 

0.1742 
0.1742 

0.0392 
0.0552 

O  .  0004 
0.0162 

+o  .  0005 

+0.0003 

Determination  of  When  certain  salts  are  brought  into  association 
Adds  Liberated  with  water,  a  tendency  to  the  formation  of  more 

Hydrolysis.  ^as{c  products  and  free  acid  in  consequence  of  the 
hydrolytic  action  of  water  becomes  evident.  Such  action  goes 
on  until  an  equilibrium  is  reached  between  the  factors  and 
products  of  reaction. 

Action  between  Iodide-iodate  Mixture  and  Certain  Salts.  —  In 
the  presence  of  the  iodide-iodate  mixture,  free  acid,  a  product  of 
the  hydrolytic  action,  may  be  constantly  removed,  and  so  the 
hydrolysis  may  proceed  further,  the  iodine  set  free  in  reaction 
of  the  iodide-iodate  mixture  upon  the  free  acid  being  obviously 
a  measure  of  the  degree  of  hydrolytic  action.  Such  action  may 
proceed  to  the  complete  hydrolysis  of  the  salt  or  may  cease  at 
an  earlier  stage,  depending  upon  the  nature  of  the  salt.  The 
behavior  of  certain  salts  in  presence  of  the  iodide-iodate  mixture 
has  been  studied  experimentally  by  Moody.*  In  Moody's 
experiments,  a  suitable  amount  of  the  salt  to  be  tested  is  put 
*  Seth  E.  Moody,  Am.  Jour.  Sci.,  [4],  xx,  181;  xxii,  176,  379,  483. 


METHODS  IN  CHEMICAL  ANALYSIS 


into  the  Voit  flask  (B)  of  the  apparatus  shown  in  Fig.  3,*  and 
10  cm.3  of  an  iodide-iodate  mixture  (i  grm.  KI.  :  0.3  grm.  KIO3) 
are  added,  the  Drexel  flask  (C)  and  trap  are  charged  with  a 
solution  containing  3  grm.  of  potassium  iodide,  hydrogen  is 
passed  from  the  generator  through  the  apparatus,  and  the  con- 
tents of  the  Voit  flask  are  heated  until  (according  to  circum- 
stances) all  or  nearly  all  the  liberated  iodine  is  collected  in  the 
Drexel  flask.  The  free  iodine,  whether  in  the  receiver  or  re- 
maining in  small  amount  in  the.  distilling  flask,  is  titrated  with 
sodium  thiosulphate. 

Aluminium  Sulphate.  —  Proceeding  in  this  manner,  Moody 
found  that  aluminium  sulphate,  though  only  partially  hydrolyzed 
at  ordinary  temperatures,  is  completely  broken  up  by  heating, 
according  to  the  reaction 

A12(S04)3  +  5  KI  +  KI03  +  3H20  =  2  A1(OH)3  +  3  K2SO4  +  312, 

and  that  aluminium  chloride  behaves  similarly. 

The  following  results  were  obtained  with  potassium  alum,  the 
potassium  sulphate  not  being  susceptible  to  hydrolytic  action. 


Approx. 
n/io 
aluminium 
potassium 
alum. 

KIO,. 

KI. 

Time  in 
minutes. 

Approx. 

M/IO 

Na2S2O,. 

A1203 
calculated 
from  iodine 
found. 

A1203 
precipi- 
tated and 
weighed. 

Difference. 

cm.s 

grm. 

grm* 

cm.8 

grm. 

grm. 

grm. 

2S 

o-3 

.0 

3° 

24-55 

0.0410 

0.0414 

—0.0004 

25 

3-3 

.0 

3° 

24.60 

0.0411 

0.0416 

—  0.0005 

25 

o-3 

.0 

25 

24.50 

o  .  0409 

0.0414 

—  0.0005 

25 

o-3 

.0 

30 

24.70 

0.0413 

0.0416 

—  0.0003 

25 

0.3 

.0 

35 

24.50 

o  .  0409 

0.0415 

—  0.0006 

25 

o-3 

.0 

30 

24-55 

0.0410 

0.0415 

—  0.0005 

25 

o-3 

.0 

25 

24.50 

0.0409 

0.0415 

—  0.0006 

Upon  the  presumption  that  the  aluminium  salt  taken  is  per- 
fectly neutral,  the  reaction  affords  means  for  determining  alu- 
minium iodometrically. 

Upon  repeating  the  experiment  with  similar  amounts  of  an 
ammonium  alum  it  was  found  that  the  amounts  of  iodine  liberated 
were  in  the  average  somewhat  excessive,  corresponding  to  0.0006 
grm.  of  A12O3  more  than  the  theory  called  for.  The  process  is, 
therefore,  less  exact  in  the  presence  of  ammonium  salts.  In  fact, 

*  See  page  4. 


APPLIANCES  AND   GENERAL  PROCEDURE 


as  will  be  shown,  ammonium  sulphate  in  the  amounts  taken  may 
be  completely  hydrolyzed  in  the  course  of  three  hours. 

Chromic  Sulphate,  Tin  Chloride.  —  Chromic  sulphate,  taken  in 
the  form  of  chrome  alum,  undergoes  complete  hydrolysis  in 
presence  of  the  iodide-iodate  mixture  with  precipitation  of 
chromic  hydroxide,  as  does  the  double  potassium  tin  chloride, 
SnCl4.2KCl. 

Iron  Sulphate.  —  Ferric  sulphate  reacts  like  aluminium  sul- 
phate according  to  the  equation 

Fe2(S04)3  +  5  KI  +  KI03  +  3  H2O  =  2  Fe(OH)»  +  3  K2SO4  +3X2. 

The  hydrolysis  of  ferrous  sulphate  is  accompanied  by  oxidation 
of  the  ferrous  hydroxide  at  the  expense  of  the  iodate,  as  follows : 

3  FeS04  +  5  KI  +  KIO3  +  3  H2O  =  3  Fe(OH)2  +  3  K2SO4  +  3  I2. 
6  Fe(OH)2  +  KI03  +  3  H2O  =  6  Fe(OH)3  +  KI. 

From  the  experimental  tests  it  appears,  however,  that  the 
hydrolysis  of  ferrous  sulphate  is  complete  in  the  presence  of  the 
iodide-iodate  mixture,  and  that  the  iodine  set  free  is  an  exact 
measure  of  the  SO4  ion  present  and  of  the  iron  in  the  ferrous  sul- 
phate of  ideal  composition. 


Iodine 

Volume. 

KI. 

KIO3. 

Time  in 
minutes. 

Approx.  «/io 
Na2S2O3. 

Iodine 
found. 

value  of 
FeS04 

Difference. 

taken. 

cm.3 

grm. 

grm. 

cm.* 

gr  111. 

grm. 

.gnu. 

40 

I  .O 

0-45 

30 

26.67 

0-3324 

0.3322 

+0.0002 

40 

I  .O 

0-45 

30 

26.68 

0.3325 

0.3322 

+  0.0003 

40 

1.0 

0-45 

30 

26.65 

0.3321 

0.3322 

—  o.oooi 

40 

1.0 

0-45 

30 

26.67 

0.3324 

0.3322 

+O.OOO2 

40 

1.0 

0.45 

3° 

26.66 

0.3323 

0.3322 

+  O.OOOI 

Cobalt  Sulphate. — Cobalt  sulphate  when  similarly  hydrolyzed 
is  likewise  oxidized  at  the  expense  of  the  iodate,  the  reactions 
following  similar  lines. 

3  CoS04  +  5  KI  +  KI03  +  3  H20  =  3  Co(OH)2  +  3  K2SO4  +  3  I2. 
6  Co(OH)2  +  KIO3  +  3  H2O  =  6  Co(OH)3  +  KI. 

The  iodine  value  obtained  in  test  experiments,  upon  the  hypoth- 
esis that  cobaltous  sulphate  is  completely  hydrolyzed,  is  closely 


64 


METHODS  IN   CHEMICAL  ANALYSIS 


comparable  with  the  iodine  equivalent  of  the  cobalt  found  by 
the  electrolytic  deposition  of  the  metal. 


Volume. 

KI. 

KIO3. 

Time  in 
hours. 

Approx. 

M/IO 

Iodine 
value  found. 

Iodine 
value  of 
CoS04 

Difference. 

812       3- 

taken. 

cm.* 

gnu  . 

grm. 

cm.3 

grm. 

grm. 

grm. 

40 

.0 

o.45 

4 

17.80 

0.2244 

0.2242 

+  O.0002 

40 

.0 

0-45 

32 

17.78 

0.2242 

0.2242 

o  .  oooo 

40 

.0 

0.45 

3f 

17-75 

0.2238 

o.  2242 

—0.0004 

40 

.0 

0-45 

4 

17.79 

0.2243 

0.2242 

+  0.0001 

40 

.0 

0-45 

4 

17.79 

0.2243 

0.2242 

+O.OOOI 

40 

.0 

0.45 

4 

17.78 

o.  2242 

o.  2242 

o  .  oooo 

Nickel  Sulphate.  —  Nickel  sulphate,  like  cobalt  sulphate,  is 
hydrolyzed  completely,  after  a  considerable  time,  in  the  presence 
of  the  iodide-iodate  mixture,'  likewise  yielding  iodine,  which  may 
be  collected  similarly  and  estimated  as  a  measure  of  the  nickel 
present.  Nickelous  hydroxide  formed  in  the  reaction  remains, 
however,  unoxidized  by  potassium  iodate  in  neutral  solution, 
and  therefore  the  following  equation  shows  the  final  products : 

3  NiS04  +  5  KI  +  KI03  +  3  H20  =  3  Ni(OH)2  +  3  K2SO4  +  3  ^* 


Volume. 

KI. 

KIO3. 

Time  in 
hours. 

Approx. 

K/IO 

Na2S203. 

Iodine 
found. 

Iodine 
value  of 
NiS04 
taken. 

Difference. 

cm.s 

grm. 

grm. 

cm.8 

grm. 

grm. 

grm. 

40 

.O 

0-45 

3 

17.87 

0.2254 

0.2255 

—  O.OOOI 

40 

.0 

0-45 

3 

17.88 

0.2256 

0.2255 

+0.0001 

40 

.O 

0.45 

3 

17.84 

0.2250 

0.2255 

—0.0005 

40 

.O 

0-45 

3 

17.87 

0.2254 

0-2255 

—  O.OOOI 

40 

.O 

0.45 

3 

17.83 

0.2249 

0.2255 

—  0.0006 

Thus  it  appears  that  nickel  sulphate  may  be  completely  hydro- 
lyzed in  the  presence  of  the  iodide-iodate  mixture,  and  that  the 
nickel  of  nickel  sulphate  of  ideal  composition  can  be  estimated 
from  the  amount  of  iodine  liberated  in  the  action  of  that  salt 
upon  the  iodide-iodate  mixture. 

Zinc  Sulphate.  —  Like  the  sulphates  of  nickel,  cobalt,  iron, 
aluminium  and  chromium,  zinc  sulphate  is  hydrolyzed  in  pres- 
ence of  the  iodide-iodate  mixture,  but  unlike  the  sulphates  of  the 
other  elements  mentioned  the  reaction  stops  short  of  complete 


APPLIANCES  AND  GENERAL   PROCEDURE 


hydrolysis,  as  shown  in  the  results  of  the  table.     The  reaction 
may  be  expressed  by  the  equation 

15  ZnSO4  +  20  KI  +  4  KIO3  +  12  H2O 

=  3  Zn5(OH)8SO4  +  12  K2SO4  +  12  I2. 


Volume. 

KI. 

KIO3. 

Time  in 
hours. 

Approx. 
n/io 
Na2S203. 

Iodine 
found. 

Equivalent 
of  SO, 
found. 

SO3  present. 

cm-3 

grm. 

grm. 

cm.3 

grm. 

grm. 

grm. 

40 

I  .O 

0-45 

* 

27-8 

0-3557 

0.1123 

0.1408 

40 

I  .0 

0-45 

i 

28.0 

0.3582 

0.1130 

0.1408 

40 

I.O 

0-45 

3 

27.8 

0-3557 

O.II23 

0.1408 

40 

1  .0 

0-45 

3 

27.8 

0-3557 

o  .  1  1  23 

0.1408 

The  mean  percentage  of  hydrolysis  here  found  is  79.90.  It 
appears  that  a  basic  sulphate  containing  5  of  Zn  to  one  of  SO4  is 
formed,  and  so  definitely  that  from  the  iodine  liberated  the  zinc 
content  may  be  calculated  with  accuracy. 

Ammonium  Sulphate.  —  When  solutions  of  ammonium  sul- 
phate are  subjected  to  heat  the  salt  is  hydrolyzed ,  and  as  the  acid 
is  increased,  either  as  a  direct  product  of  this  hydrolysis  or  by 
addition,  further  dissociation  is  inhibited.  The  decrease  in  dis- 
sociation is  dependent  upon  the  increase  of  the  acid,  and  when 
sufficient  acid  is  present  further  hydrolysis  is  entirely  prevented. 
The  amount  of  hydrolysis  is,  however,  small  under  the  most 
favorable  conditions.* 

Moody  has  studied  the  effect  of  the  iodide-iodate  mixture  in 
eliminating  the  acid  as  it  is  produced.  In  experiments  made  in 
the  manner  described  above  it  was  found  to  be  impossible  to 
collect  the  iodine  in  the  Drexel  flask  used  as  a  receiver  when 
charged  with  potassium  iodide  only,  although  it  was  evident 
that  much  iodine  came  over.  It  appeared,  upon  investigation, 
that  ammonium  iodide  and  ammonium  iodate  were  formed  by 
reaction  in  the  receiver  between  the  liberated  iodine  and  the 
ammonia  also  volatilized,  and  to  obviate  the  difficulty  sulphuric 
acid  was  added  to  the  contents  of  the  receiver  into  which  the  dis- 
tillate was  passed.  Under  these  conditions  iodine  is  obtained 
in  amount  corresponding  to  that  which  should  be  eliminated 
when  the  ammonium  sulphate  is  entirely  hydrolyzed. 
*  Bruck,  Dissertation,  Giessen,  1903. 


66 


METHODS   IN  CHEMICAL  ANALYSIS 


This  is  shown  in  the  subjoined  table: 


Volume. 

KI. 

KIO3. 

Time  in 
hours. 

H2S04 
(i*D 

in  the 
receiver. 

Approx. 

K/IO 

Na,SA. 

Iodine 
found. 

Iodine 
value  of 
(NH4),S04 
taken. 

Difference. 

cm.* 

grm. 

grm. 

cm.3 

cm.3 

grm. 

grm. 

grm. 

35 

.0 

0.30 

3 

40 

38.25 

0.4769 

0-4773 

—  0.0004 

35 

.0 

0.30 

3 

40 

38.25 

0.4769 

0-4773 

—  0.0004 

35 

.0 

0.30 

3 

40 

38.30 

°-4775 

0-4773 

+O  .  OOO2 

35 

,o 

0.30 

3 

40 

38.25 

0.4769 

0-4773 

—  O.OOO4 

35 

.0 

0.30 

3 

40 

38.23 

0.4766 

0-4773 

—  O.O007 

In  another  series  of  experiments,  the  apparatus  of  Fig.  4*  was 
used.  In  these  experiments  the  mixture  was  boiled  in  the  first 
Voit  flask  and  the  distillate  passed  from  the  first  flask,  V1,  through 
the  second  flask,  V2,  containing  50  cm.3  of  n/io  H2SO4  to  take  up 
the  ammonia,  and  then  into  the  receiver  containing  potassium 
iodide  without  acid.  The  sulphuric  acid  remaining  free  in  the 
second  Voit  flask  at  the  end  of  the  operation  was  determined  by 
titration  of  the  iodine  liberated  upon  the  addition  of  theiodide- 
iodate  mixture,  the  difference  between  this  amount  of  iodine  and 
the  iodine  equivalent  of  the  sulphuric  acid  used  being  the  meas- 
ure of  the  ammonia  absorbed.  The  iodine  passing  to  the  receiver 
was  determined  as  usual  by  titration  with  sodium  thiosulphate. 

The  results  of  these  experiments,  given  below,  show  that  the 
sulphuric  acid  neutralized  in  the  Voit  flask  is  a  measure  of  the 
ammonia,  while  the  iodine  in  the  Drexel  flask  corresponds  to  the 
sulphuric  acid  of  the  ammonium  sulphate. 

Similar  results  were  obtained  with  ammonium  chloride. 

Volume,  45  cm.3;  KI,  i  grm.;  KI03,  0.6  grm.;  Time,  j  to  3%  hours. 


Iodine  equivalent  of  ammonia 
absorbed  in  Voit  flask. 

Iodine  estimated  in  Drexel  flask. 

Iodine  value  of 

(NH4)2S04 
taken. 

Approx.  w/io 
Na2S203. 

I. 

Difference. 

Approx.  H/IO 
Na2SjO,. 

I. 

Difference. 

gnu. 

cm.* 

grm. 

grm. 

cm.* 

grrn. 

grm. 

0-4773 

38.15 

0-4757 

—  0.0016 

38.23 

0.4767 

—  0.0006 

0-4773 

38.20 

0.4763 

—  O.OOIO 

38.25 

0-4769 

—  0.0004 

0-4773 

38.15 

0-4757 

—  0.0016 

38.20 

0.4763 

—  O.OOIO 

0-4773 

38.20 

0.4763 

—  O.OOIO 

38.27 

0.4771 

—  O.OOO2 

0-4773 

38.17 

0-4759 

—0.0014 

38.20 

0.4763 

—  O.OOIO 

0-4773 

38.15 

0-4757 

—  0.0016 

38.20 

0.4763 

—  O.OOIO 

0-4773 

38.20 

0.4763 

—  O.OOIO 

38.25 

0.4769 

—  0.0004 

*  See  page  5. 


APPLIANCES  AND   GENERAL  PROCEDURE  67 

This  procedure  is  not  presented  as  an  analytical  method  for 
determining  ammonia  or  the  acid-ion,  but  to  show  that  the  effects 
of  hydrolysis  must  not  be  ignored  when  ammonium  salts  are 
heated  in  solution  with  the  iodide-iodate  mixture. 
Alums-  Basic  Moody  has  shown  *  how  the  phenomena  of  hy- 
Aiumina  and  drolysis  may  be  applied  to  the  determination  of  basic 
alumina  and  free  acid  in  the  analysis  of  alums  and 
commercial  aluminium  sulphates,  which,  beside  aluminium  sul- 
phate, may  contain  ferrous  sulphate,  ferric  sulphate,  and  zinc 
sulphate  as  impurities.  Potassium  sulphate  and  sodium  sul- 
phate, if  present,  do  not  set  free  iodine  from  the  boiling  solution 
containing  the  iodide-iodate  mixture.  The  determinations  of  the 
ferrous  iron,  the  ferric  iron,  the  zinc,  and  the  ammonia  furnish 
data  from  which  the  equivalent  amounts  of  sulphuric  acid,  to  be 
taken  into  account  in  the  reckoning  of  the  free  acid  or  basic 
alumina,  may  be  calculated.  The  behavior  of  these  commercial 
products  toward  the  iodide-iodate  mixture  affords,  therefore,  an 
easy  method  of  determining  basic  alumina  or  free  acid,  as  the 
case  may  be. 

Following  are  the  details  of  treatment  : 

I.  A  sample  of  15  grm.  is  weighed  and  treated  with  water. 
The  solution  is  filtered  and  made  up  to  I  liter.     The  material 
which  does  not  dissolve  is  dried  at  1 00°  and  weighed  as  insoluble 
material. 

II.  Of  the  solution,  a  portion  of  25  cm.3  is  titrated  directly 
with  standard  potassium  permanganate  to  find  the  amount  of 
iron  in  the  ferrous  salt,  and  from  this  is  calculated  the  ferrous 
oxide. 

III.  Of  the  solution,  another  portion  of  25  cm.3  is  treated  with 
zinc  to  reduce  the  ferric  salt  and  then  titrated  with  permanganate 
to  give  the  total  iron.     From  the  difference  between  the  total 
iron  and  the  ferrous  iron  is  calculated  the  ferric  oxide. 

IV.  Of  the  solution,  a  portion  of  25  cm.3  is  diluted  to  50  cm.8, 
treated  with  3  grm.  of  sodium  acetate  and  I  cm.3  of  acetic  acid, 
and  electrolyzed  with  the  use  of  the  rotating  cathode  f  by  a 
current  of  about  2  amperes  for  30  minutes.     The  deposit  of  zinc, 
including  some  iron,  is  washed  with  alcohol,  dried  and  weighed. 

*  Am.  Jour.  Sci.,  [4],  xxii,  483. 
t  See  page  n. 


68  METHODS  IN   CHEMICAL  ANALYSIS 

The  solution  of  the  deposit  in  sulphuric  acid  is  titrated  with 
permanganate,  and  the  amount  of  iron  thus  found  is  deducted 
from  the  total  weight  of  the  deposit  to  give  the  amount  of  zinc. 
From  the  zinc  is  calculated  the  zinc  oxide. 

V.  Of  the  solution,  a  portion  of  25  cm.3  is  drawn  from  a  bu- 
rette into  the  Voit  flask  of  the  distillation  apparatus,  a  solution 
(10  cm.3)  containing  0.3  grm.  of  potassium  iodate  and  I  grm.  of 
potassium  iodide  is  added,  the  mixture  boiled,  and  the  iodine, 
collected  in  the  receiver  charged  with  water  containing  3  grm. 
of  potassium  iodide  (and  acidified  with  sulphuric  acid  in  case  the 
substance  contains  ammonia),  is  titrated  with  sodium  thiosul- 
phate. 

The  iodine  set  free  corresponds  to  the  various  oxides,  to  am- 
monia and  to  sulphuric  acid  in  the  following  proportions : 


A1203 

61; 

Fe2O3 

61; 

FeO 

a  Is 

5ZnO 

81; 

NH3 

I; 

H2SO4 

a  Is 

it  is  the  total  iodine. 

VI.  Of  the  solution,  a  portion  of  25  cm.3  is  treated  in  an  open 
beaker  with  I  grm.  of  potassium  iodide  and  0.3  grm.  of  potassium 
iodate,  the  mixture  is  boiled  until  nearly  all  free  iodine  is  ex- 
pelled, the  precipitate  is  filtered  on  paper,  ignited  carefully,  and 
weighed  as  A12O3,  Fe2O3  and  ZnO,  the  total  oxides. 

Given  the  total  iodine  liberated  in  the  distillation  process,  the 
weight  of  the  total  oxides  obtained  in  the  parallel  boiling  process, 
the  ferrous  oxide  and  ferric  oxide  by  the  permanganate  titrations 
and  the  zinc  oxide  deduced  from  the  corrected  electrolytic  deter- 
mination, the  total  alumina  and  the  basic  alumina  or  free  acid  (as 
the  case  may  be)  are  easily  calculated. 

Total  oxides  —  (ferric  oxide  +  ferrous  oxide  +  zinc  oxide)  = 
total  alumina. 

/6X  I2^97\Qr  (7454)  x  total  alumina  = 

iodine  equivalent  to  total  alumina. 


APPLIANCES  AND   GENERAL  PROCEDURE 


69 


or  (4.767)  X  ferric  oxide  = 

iodine  liberated  by  ferric  sulphate. 
/2Xi26.97\  or  ( 

iodine  liberated  by  ferrous  sulphate. 

/8  X  l26-97\  or  (2496)  x  zinc  oxide  = 
4  / 


5  X  814 


/  --  i 


iodine  liberated  by  zinc  sulphate. 
(7.469)  X  ammonia  = 
iodine  liberated  by  ammonium  sulphate. 


Total  iodine  —  (iodine  corresponding  to  total  alumina,  ferric 
sulphate,  ferrous  sulphate,  zinc  sulphate,  ammonium  sulphate)  = 
differential  iodine. 

Differential  iodine  (if  positive) 


Differential  iodine  (if  negative) 


X 


=  basic  alumina" 


The  results  of  analyses  of  four  specimens  of  alums  are  given 
in  the  following  table : 

Percentage  Composition. 


AUOs. 

FeO.* 

ZnO. 

NHs. 

Insoluble 
material. 

Total. 

Combined. 

Basic. 

No.  I: 

(i) 

14.48 

J3-49 

0-99 

0-43 

3-70 

None. 

0.61 

(2) 

14.28 

I3-46 

0.82 

0.44 

3-70 

None. 

0.61 

No.  II: 

d) 

15-94 

14.21 

i-73 

0-43 

1-47 

None. 

0.21 

(2) 

I5-90 

14-34 

1-56 

0-45 

1-34 

None. 

O.2I 

No.  Ill: 

(i) 

J5-59 

14.81 

0.78 

0-34 

0.73 

None. 

0.71 

(2) 

15-97 

14.80 

1.17 

0.36 

0.82 

None. 

0.71 

D: 

(i) 

16.59 

15.24 

1-35 

0.24 

O.II 

None. 

0.61 

(2) 

16.37 

15.18 

1.19 

0.24 

0.19 

None. 

0.61 

With  a  trace  of 


70  METHODS  IN   CHEMICAL  ANALYSIS 

The  Use  of  the  Bromide-bromate  Mixture  and  the  Estimation  of 
the  Bromine  Evolved. 

The  reaction  of  a  mixture  of  potassium  bromide  and  potassium 
bromate  upon  aluminium  sulphate  has  been  studied  by  Gooch  and 
Osborne*  with  the  aid  of  the  apparatus  previously  described. f 
Assuming  that  the  acidic  ion  is  entirely  liberated  from  the 
aluminium  salt,  the  reaction  should  follow  the  equation 

2  KA1(S04)2  +  5  KBr  +  KBrO3  +  3  H2O  =  2  Al(OH), 
+  4  K2S04  +  3  Br2. 

The  precipitation  proves  to  be  complete,  or  nearly  so;  but  the 
process  of  hydrolysis  does  not  go  easily  to  the  point  of  forming 
aluminium  hydroxide.  A  reasonable  excess  of  the  bromide-bro- 
mate  mixture  is  able  in  a  moderate  time  to  carry  the  hydrolysis 
of  aluminium  sulphate  to  a  fairly  definite  point  corresponding 
nearly  to  the  removal  of  five-sixths  of  the  acidic  ion,  while  the 
iodide-iodate  mixture  under  similar  conditions  of  action  removes 
practically  all  the  acidic  ion. 

With  a  very  large  increase  in  the  concentration  of  the  bro- 
mide and  bromate  and  prolonged  boiling,  the  completion  of  the 
hydrolysis  to  the  point  of  liberating  bromine  equivalent  to  the 
entire  amount  of  the  acidic  ion  is  very  nearly  realized. 

Like  the  iodide-iodate  mixture,  the  bromide-bromate  mix- 
ture is  a  very  delicate  indicator  of  free  acid;  0.00018  grm.  of 
sulphuric  acid  proving  to  be  sufficient  to  liberate  bromine  from 
the  bromide-bromate  mixture  when  boiled  in  the  Voit  flask  under 
the  experimental  conditions. 

Some  experiments  to  test  the  effect  of  a  mixture  of  potassium 
chloride  and  potassium  chlorate  upon  aluminium  sulphate  indi- 
cated that  the  hydrolysis  under  otherwise  similar  conditions  is 
very  slight  compared  with  that  produced  by  the  bromide- 
bromate  mixture  or  by  the  iodide-iodate  mixture. 

The  Reaction  of  Iodine  with  Alkali  Hydroxides. 

When  the  solution  of  a  metallic  hydroxide  is  acted  upon  by 
iodine  at  a  temperature  high  enough  to  decompose  the  small 
amounts  of  hypoiodites  that  might  otherwise  be  present,  the  final 

*  F.  A.  Gooch  and  R.  W.  Osborne,  Am.  Jour.  Sci.,  [4],  xxiv,  167. 
t  See  pages  4,  61. 


APPLIANCES  AND  GENERAL  PROCEDURE        71 

action  results  in  the  formation  of  an  exactly  neutral  mixture  of 
iodate  and  iodide  according  to  the  equation 

6  KOH  +  3  I2  =  RI03  +  5  RI  +  3  H2O. 

In  a  process  described  elsewhere*  for  the  determination  of 
carbon  dioxide,  Phelps  f  has  applied  this  reaction  to  the  deter- 
mination of  barium  hydroxide.  In  this  process,  standard  iodine 
is  made  to  act  in  a  suitable  apparatus  upon  the  barium  hydroxide 
and  the  excess  of  it  is  determined. 

Instead  of  determining,  according  to  the  procedure  of  Phelps, 
the  amount  of  iodine  left  over  in  the  action  of  an  excess  of  that 
element  upon  the  alkali  hydroxide,  Walker  and  Gillespiet  pre- 
fer to  expel,  by  boiling,  all  free  iodine  remaining  after  action 
and  to  determine  the  amount  of  iodine  liberated  by  acting 
with  an  acid  from  the  residual  mixture  of  iodate  and  iodide. 
According  to  the  procedure  recommended,  an  excess  of  deci- 
normal  iodine  is  drawn  into  an  Erlenmeyer  beaker  and  the  de- 
sired amount  of  alkali  hydroxide  is  run  in  rapidly.  The  beaker, 
closed  by  a  little  trap,§  made  of  a  calcium  chloride  drying  tube, 
to  prevent  appreciable  loss  by  spattering,  is  placed  over  a  low 
flame,  and  the  contents  boiled  until  the  last  trace  of  the  excess  of 
iodine  has  been  volatilized  from  the  solution  and  the  trap.  The 
volume  is  carefully  regulated  before  and  during  the  boiling,  being 
kept  as  small  as  possible,  usually  amounting  to  about  100  cm.3  at 
the  start  and  35  cm.3  at  the  close.  In  the  case  of  barium  hy- 
droxide, care  has  to  be  taken  to  keep  the  dilution  sufficient  to 
prevent  the  separation  of  the  crystalline  barium  iodate,  which  is 
soluble  with  difficulty.  To  steady  the  ebullition  a  little  spiral  of 
platinum  is  introduced  into  the  beaker.  After  the  boiling,  the 
colorless  solution,  containing  a  neutral  mixture  of  iodate  and 
iodide,  is  cooled  in  running  water,  and  treated  with  10  cm.3  of 
dilute  acid  — in  the  case  of  barium  hydroxide,  dilute  [i :  3]  hydro- 
chloric acid,  to  save  the  inconvenience  of  working  in  the  pres- 
ence of  precipitated  barium  sulphate;  with  potassium  hydrox- 
ide, dilute  [i :  3]  sulphuric  acid.  The  liberated  iodine  is  titrated 

*  See  page  231. 
f  Am.  Jour.  Sci.,  [4],  ii,  70. 

$  Claude  F.  Walker  and  David  H.  M.  Gillespie,  Am.  Jour.  Sci.,  [4],  vi, 
455- 

§  See  Fig.  6. 


72 


METHODS  IN  CHEMICAL  ANALYSIS 


directly  with  sodium  thiosulphate,  in  the  presence  of  starch. 
Tests  of  this  method  in  comparison  with  that  of  Phelps  showed 
a  fair  agreement,  though  the  results  of  both  methods  were  in- 
variably lower  by  a  small,  nearly  constant  amount  than  those  ob- 
tained by  gravimetric  estimations  and  by  the  Groger  process.* 
This  error  of  the  Phelps  process  and  its  modification  is  probably 
due  to  the  action  of  atmospheric  carbon  dioxide  on  the  hydroxide 
solution  during  the  short  time  it  is  exposed.  While  it  affects 
considerably  the  value  of  the  method  as  a  means  of  accurately 
determining  the  absolute  amount  of  hydroxide  present  in  a  given 
volume  of  solution,  it  apparently  does  not  so  seriously  affect  the 
accuracy  of  the  differential  method  founded  on  the  original  Phelps 
process  or  its  modification. 

Analyses  ofn/io  Hydrochloric  Acid  Solution. 

(By  adding  to  excess  of  w/io  Ba(OH)2,  boiling  with  excess  of  iodine  to 
colorlessness,  and  acidifying  the  residue.) 


Ba(OH)2 

HClby 

HC1  taken. 

Ba(OH)2  taken. 

neutralized  by 

HC1  found. 

Groger 

Variation. 

HC1. 

method. 

cm.* 

grm. 

grm. 

grm. 

gnu  . 

grm. 

IS 

0.17 

0.1128 

o  .  0480 

0.0481 

+O.OOOI 

IS 

0.17 

0.1118 

0.0475 

0.0481 

—0.0006 

IS 

0.17 

O.III2 

0.0473 

0.0481 

—0.0008 

25 

o.  26 

0.1860 

0.0791 

0.0801 

—  o.ooio 

25 

0.26 

0.1866 

0.0794 

0.0801 

—0.0007 

35 

0.34 

0.2634 

O.  1  1  2O 

O.  II2I 

—  o.oooi 

35 

0.34 

0.2603 

0.1107 

O.II2I 

—0.0014 

Analyses  ofn/io  Hydrochloric  Acid  Solution. 

(By  adding  to  excess  of  w/io  KOH,  boiling  with  excess  of  iodine  to  color- 
lessness, and  acidifying  the  residue.) 


KOH 

HClby 

HC1  taken. 

KOH  taken. 

neutralized 

HC1  found. 

Gioger 

Variation. 

by  HC1. 

method. 

cm.s 

grm. 

grm* 

grm. 

grm. 

grm. 

2O 

0.14 

0.0972 

0.0633 

O  .  0641 

—  0.0008 

2O 

0.14 

0.0975 

0  .  0634 

O  .  0641 

—  0.0007 

25 

0.14 

O.I222 

0.0795 

O  .  0801 

—  O.OOO6 

25 

0.14 

O.I2O7 

0.0785 

o  .  0801 

—  O.OOI6 

In  applying  the  Walker-Gillespie  modification  to  the  indirect 
determination  of  hydrochloric  acid  and  sulphuric  acid,  the  acid 

*  See  page  59. 


APPLIANCES  AND   GENERAL  PROCEDURE 


73 


solution  to  be  analyzed  is  drawn  into  an  Erlenmeyer  beaker,  an 
excess  of  decinormal  iodine  is  added,  and  a  measured  excess  of 
standardized  alkali  is  introduced.  The  beaker  is  trapped,  the 
solution  boiled,  and  the  residue  cooled,  acidulated,  and  titrated 
with  thiosulphate  in  presence  of  starch.  Test  results  of  the 
process  are  given  in  the  accompanying  tables. 

Analyses  ofn/io  Sulphuric  Acid  Solution. 

(By  adding  to  excess  of  n/io  Ba(OH)2,  boiling  with  excess  of  iodine  to 
decoloration,  and  acidifying  the  residue.) 


H2SO< 

taken. 

Ba(OH)2 
taken. 

Ba(OH)2 
neutralized 
by  H,SO«. 

H,S04 

found. 

H2SO<  by 
Groger  method. 

Variation. 

cm.3 

grin. 

grm. 

grm. 

grm. 

grm. 

(i)     10 

O.  21 

o  .  0884 

O  .  0506 

o  .  0496 

+O.OOIO 

(2)       10 

O.  21 

0.0880 

0  .  0503 

o  .  0496 

+0.0007 

(3)    is 

0.30 

0.1328 

0-0754 

0-0745 

+0.0009 

(4)     IS 

0.30 

0.1313 

0.0751 

0-0745 

+0.0006 

(5)     25 

0-43 

0.2168 

0.1239 

o  .  i  240 

—  o.oooi 

(6)     30 

0-43 

o  .  2600 

o.  1481 

0.1489 

—0.0008 

The  reaction  between  iodine  and  the  hydroxides  of  potassium 
and  barium  in  hot  solution  is,  therefore,  regular  and  complete 
under  analytical  conditions,  and  not  appreciably  affected  by  the 
mass  action  of  considerable  excesses  of  iodine.  It  is  best  applied  in 
analysis  by  treating  the  alkali  with  an  excess  of  iodine,  removing 
this  excess  by  boiling,  and  estimating  the  iodine  in  the  residue. 
While  mechanical  difficulties  and  the  interfering  action  of  carbon 
dioxide  may  affect  the  extreme  accuracy  of  the  process  as  a 
direct  means  for  analyzing  alkalies,  it  may  be  indirectly  applied 
with  fair  accuracy  to  the  analysis  of  various  acids  and  possibly 
to  other  compounds.  The  reaction  between  iodine  and  alkali 
carbonates,  on  the  contrary,  is  irregular  and  cannot  be  made  the 
basis  of  an  analytical  process. 


CHAPTER   II. 
THE   ALKALI   METALS. 

SODIUM. 
The  Detection  of  Sodium. 

THE  application  of  the  spectroscopic  method  to  the  detection 
of  potassium  and  sodium  in  ordinary  analysis  is  unsatisfactory 
on  account  of  its  failure,  except  in  delicate  quantitative  compari- 
sons, to  give  any  idea  as  to  the  quantity  of  either  element  indi- 
cated; and,  since  the  most  minute  quantity  of  either  element  is 
sufficient  to  produce  its  characteristic  line  in  the  spectroscope, 
and  many  of  the  reagents  employed  in  analysis  contain  a  trace  of 
alkali,  the  spectroscopic  indication  is  often  misleading.  While 
to  the  careful  observer  the  presence  or  absence  of  potassium  in 
appreciable  amount  is  revealed,  the  evidence  as  to  the  quantity 
of  the  ubiquitous  element  sodium  is  practically  worthless. 

Kreider  and  Breckenridge  *  have  developed  a  method  for  the 
detection  of  sodium  based  upon  the  utilization  of  the  perchlorate 
method  for  the  preliminary  separation  of  potassium.  The  in- 
solubility of  potassium  perchlorate  and  the  easy  solubility  of 
sodium  perchlorate  in  97  per  cent  alcohol  afford  means  for  the 
separation  of  these  elements  as  well  as  for  the  identification  of 
the  former.f  By  converting  to  the  chloride  the  sodium  per- 
chlorate in  the  alcoholic  filtrate  from  the  precipitated  potassium 
salt,  exceedingly  small  amounts  of  sodium  may  be  detected. 
For  this  purpose  the  passing  of  gaseous  hydrochloric  acid  into 
the  alcoholic  solution  of  the  sodium  perchlorate,  kept  cool,  has 
proved  most  effectual,  the  dehydrating  effect  of  the  acid  upon 
the  alcohol  greatly  increasing  the  insolubility  of  the  sodium 
chloride.  The  delicacy  of  this  test  for  sodium  is  shown  in  the 
results  of  the  table. 

By  the  use  of  10  cm.3  of  97  per  cent  alcohol  with  gaseous 
hydrochloric  acid  0.0003  grm-  °f  sodium  oxide  can  be  found  with. 

*  D.  Albert  Kreider  and  J.  E.  Breckenridge,  Am.  Jour.  Sci.,  [4],  ii,  263. 
t  See  page  88. 

74 


THE  ALKALI  METALS 


75 


certainty,  and  when  the  alcohol  is  allowed  to  become  saturated 
with  the  gas  even  0.00006  grm.  The  quantity  of  alcohol,  10  cm.3, 
is  sufficient  for  all  purposes,  since  this  amount  will  dissolve  about 
2  grm.  of  sodium  perchlorate ;  but  even  in  40  cm.3  0.0002  grm. 
of  sodium  oxide  may  be  seen  distinctly.  It  is  evident  that  this 
method  may  also  be  applied  to  the  quantitative  determination 

of  sodium. 

Test  for  Sodium. 


NaClO4  taken, 
grm. 

Na2O  equivalent, 
grm. 

97  per  cent  alcohol. 
cm.s 

Indication. 

O.OIOO 

0.00250 

IO 

Very  strong. 

0.0050 

O.OOI25 

IO 

Strong. 

0.0040 

O.OOIOO 

IO 

Strong. 

0.0030 

0.00075 

IO 

Strong. 

0.0030 

0.00075 

10 

Good. 

O.OO2O 

o  .  00050 

10 

Good. 

O.OO2O 

o  .  00050 

10 

Good. 

O.OOIO 

0.00025 

IO 

Good. 

0.0005 

O.OOOI2 

IO 

Trace. 

0.0003 

o  .  00006 

IO 

Trace. 

0.0001 

o  .  00003 

IO 

None. 

0.0000 

0  .  00000 

IO 

None. 

O.OOIO 

0.00025 

40 

Distinct. 

In  the  following  table  are  recorded  the  results  of  experiments 
in  which  the  perchlorates  of  sodium  and  potassium  were  sepa- 
rated by  97  per  cent  alcohol  and  the  sodium  test  made  upon  the 
alcoholic  solution.  The  sodium  and  potassium  salts  were  drawn 
from  separate  standard  solutions  of  the  purified  perchlorates. 

Separation  of  Potassium  with  Test  for  Sodium. 


KC104 

taken. 

grrn. 

K2O 

equivalent. 

grm. 

NaC104 
taken. 

grm. 

Na20 
equivalent. 

grm. 

Indication  for 
potassium. 

Indication  for 
sodium. 

0.0500 

0.01699 

0.0500 

0.01250 

Strong. 

Strong. 

O.O2OO 

0.00680 

O.O2OO 

0.00500 

Strong. 

Strong. 

O.OIOO 

0.00340 

O.OIOO 

0.00250 

Strong. 

Strong. 

0.0050 

0.00170 

0.0050 

0.00125 

Strong. 

Strong. 

O.OO4O 

0.00136 

o  .  0040 

O.OOIOO 

Good. 

Good. 

O.OO3O 

O.OOIO2 

O.OO3O 

0.00075 

Good. 

Good. 

O.OO2O 

O.OOO68 

O.OO2O 

0.0005 

Good. 

Good. 

O.OOIO 

0.00034 

O.OOIO 

0.00025 

Good. 

Good. 

0.0005 

0.00017 

0.0005 

O.OOOI2 

Trace. 

Trace. 

0.0003 

O.OOOIO 

0.0003 

O.OOOO7 

Trace. 

Trace. 

O  .  OOOI 

o  .  00003 

0.0001 

O.OOOO3 

Faintest  trace. 

None. 

o  .  oooo 

o.ooooo 

O.OIOO 

O.OO25O 

None. 

Strong. 

O.OIOO 

0.00340 

o.oooo 

0.00000 

Strong. 

None. 

76  METHODS  IN  CHEMICAL  ANALYSIS 

After  evaporating  to  dryness  on  the  steam  bath,  the  residue  was 
treated  with  10  cm.3  of  97  per  cent  alcohol,  the  insoluble  potas- 
sium perchlorate  was  removed  by  filtering  through  a  dry  paper 
filter  and  dry  funnel  into  a  dry  test  tube,  and  the  filtrate  saturated 
with  gaseous  hydrochloric  acid. 

The  results  show  that  sodium  and  potassium  perchlorates  when 
associated  in  any  proportion  may  be  separated  by  97  per  cent 
alcohol  with  exactness,  and  that  the  sodium  may  be  indicated 
in  the  filtrate  with  great  delicacy  and  certainty  by  the  action  of 
gaseous  hydrochloric  acid. 

For  the  conversion  of  other  salts  of  sodium  and  potassium  to 
perchlorates  for  the  purpose  of  making  the  test  for  sodium,  it  is 
obvious  that  perchloric  acid  free  from  sodium  must  be  employed. 
The  perchloric  acid  prepared  according  to  the  method  described 
later,*  —  by  heating  sodium  chlorate  to  form  the  perchlorate, 
destroying  any  residual  chlorate  by  treatment  with  the  strongest 
hydrochloric  acid,  separating  sodium  chloride  by  filtration  on 
asbestos  in  the  filtering  crucible,  and  removing  the  excess  of 
hydrochloric  acid  by  evaporation,  —  while  answering  perfectly 
well  for  the  detection  of  potassium,  is  inapplicable  to  the  test  for 
sodium,  because  of  the  small  amount  of  this  element  which  the 
acid  always  contains  on  account  of  the  partial  solubility  of 
sodium  chloride  in  hydrochloric  acid.  This  perchloric  acid  must, 
therefore,  be  purified  by  distillation,  and  to  prevent  loss  by  de- 
composition the  process  must  be  carried  on  under  diminished 
pressure.  To  obviate  violent  ebullition,  only  acid  previously 
concentrated  to  the  fuming  point  should  be  subjected  to  the 
distillation  process,  and  this  only  in  small  amounts.  Rubber 
stoppers  or  connectors  are  not  to  be  used  where  the  acid  may 
condense  upon  them  and  flow  back  into  the  flask,  since  oxidizable 
matter  carried  back  causes  explosions  which  vary  in  force  and 
seriousness  according  to  circumstances. 

To  construct  a  suitable  apparatus  a  strong  side-neck  flask  is 
selected,  the  bottom  covered  to  the  depth  of  I  cm.3  with  fine 
chips  of  porcelain,  and  into  the  neck  is  sealed  a  stoppered  funnel 
reaching  well  into  the  bulb.  The  stopcock  of  this  funnel  is 
carefully  cleansed  and  lubricated  with  metaphosphoric  acid  ob- 
tained by  boiling  sirupy  orthophosphoric  acid  until  the  tem- 
perature of  350°  C.  has  been  attained.  The  side  neck  of  the 

*  Page  89. 


THE  ALKALI  METALS  77 

flask  is  inclined  upward  for  a  short  distance  before  being  bent 
into  the  receiver,  with  which  it  is  connected  by  a  rubber  stopper 
through  which  the  tube  extends  for  a  safe  distance.  An  ordinary 
bottle  of  250  cm.3  capacity  serves  for  a  receiver.  This  is  closed 
by  a  doubly  perforated  stopper,  and  through  one  of  the  per- 
forations the  adapter  from  the  condenser  is  inserted,  while 
through  the  other  connection  is  made  with  a  small  glass  bulb 
and  absorption  tube,  filled  with  stick  potash  to  take  up  any 
chlorine  evolved  in  the  inevitable  slight  decomposition  of  the 
acid,  which  in  turn  is  connected  with  an  automatic  mercury 
pump.  The  whole  flask  is  surrounded  by  a  cylinder  of  thin 
sheet  iron  closed  below,  while  the  upper  opening  is  protected 
by  an  asbestos  cover  in  order  that  the  heat  may  be  uniformly 
applied  up  to  the  point  at  which  condensed  acid  flows  into  the 
receiver. 

In  making  a  distillation,  3  cm.3  or  4  cm.3  of  the  concentrated 
acid  are  admitted  to  the  flask  through  the  stoppered  funnel, 
the  pump  is  started,  and  when  the  pressure  has  been  reduced 
to  about  8  mm.  the  heat  is  raised  to  about  130°  and  the  dis- 
tillation begins.  During  the  process  the  pressure  is  kept  at 
about  3  mm.  to  5  mm.,  and  the  acid  is  admitted  at  about  the 
same  rate  at  which  the  distilled  product  drops  from  the  con- 
denser. Under  the  conditions  described,  the  process  will  yield 
per  hour  from  25  cm.3  to  40  cm.3  of  the  dihydrate  of  perchloric 
acid,  the  most  concentrated  form  in  which  the  acid  is  stable. 
Of  this  product  o.i  grm.  of  potassium  oxide  requires  for  pre- 
cipitation 0.16  cm.3. 

With  pure  perchloric  acid  at  hand  the  separation  of  the  sodium 
salt  from  a  mixture  of  the  pure  chlorides  of  sodium  and  potassium, 
preparatory  to  making  the  sodium  test,  requires  only  treatment 
of  the  mixture  with  perchloric  acid,  evaporation  on  the  water 
bath  until  the  white  fumes  of  perchloric  acid  appear,  treatment  of 
the  residue  with  97  per  cent  alcohol,  and  filtration.  In  general, 
however,  it  is  necessary  to  remove  certain  interfering  substances 
before  applying  the  method.  While  potassium  may  be  safely 
tested  for  in  the  presence  of  other  bases  and  acids,  except  ammo- 
nium, caesium  and  rubidium,  and  sulphuric  acid,  there  are  many 
elements  the  insolubility  of  whose  chlorides  in  alcohol  necessitates 
their  removal  before  testing  for  sodium.  But  among  the 
common  alkalies  ammonium  is  the  only  one  whose  presence  is 


METHODS  IN  CHEMICAL  ANALYSIS 


objectionable.     Lithium  does  not  affect  the  test  for  either  potas- 
sium or  sodium. 

In  the  experiments  made  with  potassium  and  sodium  salts  asso- 
ciated with  salts  of  other  common  elements,  the  results  of  which 
are  recorded  below,  the  following  treatment  was  adopted.  The 
several  groups  of  bases  were  successively  removed  in  the  ordinary 
way:  Hydrogen  sulphide  in  ammoniacal  solution  removed  the 
lead,  mercury,  copper  and  zinc.  Barium  and  calcium  were  re- 
moved by  ammonium  carbonate,  the  final  filtrate  being  evapo- 
rated and  the  residue  ignited  to  volatilize  ammonium  salts.  The 
residue  was  dissolved  and  treated  with  barium  hydroxide  for  the 
removal  of  magnesium,  and,  after  filtering,  the  barium  was  again 
removed  by  ammonium  carbonate,  and  the  filtrate  evaporated. 
The  residue  was  then  ignited  as  before,  treated  with  10  cm.3  of 
boiling  water  and  the  solution  filtered  in  order  to  remove  the 
organic  matter  usually  found  at  this  stage  of  the  treatment.  To 
the  filtrate  was  added  o.i  to  0.5  cm.3  of  pure  perchloric  acid, 
about  1.7  sp.gr.,  according  to  the  amount  of  residue,  and  the 
mixture  was  evaporated  over  the  steam  bath  until  the  white 
fumes  of  perchloric  acid  appeared.  When  the  quantity  of  so- 
dium is  large  it  is  safer  to  evaporate  several  times  in  order  to 
secure  a  complete  conversion  to  the  perchlorate. 

Potassium  and  Sodium  in  Mixtures  of  Salts  of  Other  Elements. 


Pb,  Cu,  Al,  Fe,  Zn, 
Ba  Ca  and  Mg 

KjO  taken. 

Na2O  taken. 

Indication  for 

,  Indication  for 

as  nitrates. 

potassium. 

sodium. 

grm. 

gnu. 

grm. 

0.0500  of  each. 

o.oooo 

o.oooo 

Faintest  trace. 

Trace. 

0.0500  of  each. 

0.0017 

O.OOI2 

Good. 

Good. 

o.iooo  of  each. 

0.0000 

0.0000 

Faintest  trace. 

Trace. 

o.iooo  of  each. 

0.0000 

0.0005 

Faintest  trace. 

Good. 

The  fact  that  minute  traces  of  sodium  and  potassium  are 
found  in  the  blank  tests  is  to  be  expected  from  the  delicacy  of 
the  method  when  it  is  remembered  that  but  very  few  of  the 
so-called  chemically  pure  reagents  are  absolutely  free  from  so- 
dium, and  that  even  distilled  water  kept  in  glass  vessels  con- 
tains a  trace  of  the  alkali  elements.  However,  the  indication 
for  sodium  in  the  blank  tests  appeared  only  as  a  cloudiness  and 
after  complete  saturation.  When  the  quantity  of  sodium  oxide 


THE  ALKALI   METALS 


79 


present  is  not  less  than  0.0005  grm.  the  precipitate  appears  in 
granular  form  and  before  the  alcohol  is  completely  saturated. 
The  method  is  all  that  could  be  desired  for  the  qualitative  deter- 
mination of  sodium. 

The  Estimation  of  Sodium  as  the  Pyrosulphate. 

Sodium  sulphate  in  water  solution  may  be  recovered  by  simple 
evaporation  and  ignition.  When  free  sulphuric  acid  is  also 
present  the  acid  sulphate  is  first  formed,  then,  as  the  tempera- 
ture rises,  the  pyrosulphate,  and  gradually,  at  red  heat,  the 
neutral  sulphate.  In  quantitative  estimations  it  is  usual  to  weigh 
in  the  form  of  neutral  sulphate  and  to  hasten  conversion  to  this 
condition  by  making  the  final  ignition  in  the  atmosphere  pro- 
duced by  ammonium  carbonate  either  projected  into  the  hot  cru- 
cible,* or  placed  in  the  crucible  before  the  application  of  heat.f 

In  a  study  of  the  behavior  of  the  sulphates  of  the  alkali  ele- 
ments, Browning  J  has  shown  that  by  holding  the  temperature 
between  250°  and  270°  during  the  heating  of  sodium  sulphate 
with  the  excess  of  sulphuric  acid,  sodium  pyrosulphate,  Na2S2Oy, 
may  be  formed  with  a  degree  of  exactness  which  makes  it  possible 
to  estimate  sodium  as  that  salt. 

Results  of  this  procedure  are  given  in  the  table,  and,  for  com- 
parison, the  results  obtained  in  the  usual  process  of  weighing  as 
the  neutral  sulphate  after  moistening  with  ammonium  hydroxide 
and  then  igniting  strongly. 

Sodium  as  the  Pyrosulphate. 


NaCl  taken, 
grm. 

NmSA 

calculated, 
grm. 

Na,S20T 

found. 

grm. 

Error, 
grm. 

Na,SO4 
calculated. 

grm. 

NasSO4 

found. 

grm  . 

Error. 

giro. 

0.1042 
o  1028 

0.1978 
O    I(K2 

0.1972 
O    ICK2 

—  O.OOO6 
O  OOOO 

0.1266 

0.1254 

—  O.OOI2 

0.1093 
0.1402 

0.2075 
O.2662 

o  .  2065 

0.2651 

—  O.OOIO 
—  O.OOII 

0.1328 
0.1703 

0.1320 
0.1696 

—  O.OOOS 
—  0.0007 

Salts  of   potassium   also    yield    when   similarly   treated   the 
pyrosulphate, §    while    salts   of   caesium   and    rubidium    remain 

*  Fresenius,  Quant.  Anal.,  trans,  Cohn,  pages  161,  165. 
t  Treadwell,  Anal.  Chem.,  trans.  Hall  (1911),  page  42. 
t  Philip  E.  Browning,  Am.  Jour.  Sci.,  [4],  xii,  301. 
§  See  page  92. 


8o  METHODS  IN  CHEMICAL  ANALYSIS 

under  the  conditions  of  temperature  in  the  condition  of  acid 
sulphates.*  Lithium  salts  apparently  do  not  yield  the  acid 
sulphate  or  the  pyrosulphate  in  stable  form. 

POTASSIUM. 
The  Spectroscopic  Detection  and  Determination  of  Potassium. 

Bunsen  and  Kirchoff  originally  determined  the  delicacy  of  the 
spectroscopic  test  for  potassium  by  exploding  in  a  darkened 
room  a  mixture  of  potassium  chlorate  with  milk  sugar,  and 
observing  the  amount  of  finely  divided  chloride  which  it  was 
necessary  to  diffuse  through  the  given  space  in  order  to  bring 
out  unmistakably  the  spectrum  of  the  metal.  These  investiga- 
tors were  able  to  state  that  the  presence  of  no  more  than  ^^-^  of 
a  milligram  of  the  potassium  salt  is  sufficient  to  give  to  the  flame 
the  characteristic  spectrum  of  the  element.  By  similar  methods, 
the  delicacy  of  the  tests  for  lithium  carbonate  and  sodium  chlorate 
were  shown  to  be  a  thousand  times  and  three  thousand  times  as 
delicate  respectively.  The  practical  detection  of  lithium  and 
sodium  spectroscopically  is  extremely  easy  and  satisfactory,  the 
only  difficulty  being  that  the  exceeding  delicacy  of  the  sodium 
test  and  the  ubiquitousness  of  sodium  salts  often  make  a  decision 
doubtful  as  to  whether  that  element  is  present  appreciably  in  the 
substance  under  examination,  or  by  accident.  With  potassium 
the  case  is  different. 

Detection  of  Gooch  and  Hart  f  have  succeeded  in  showing  that 

Potassium.  while  the  simple  method  in  vogue  for  developing  the 
luminosity  of  lithium  and  sodium  —  the  dipping  of  a  single  loop 
of  platinum  wire  in  the  liquid  or  solid  substance,  and  the  placing 
of  the  loop  in  the  Bunsen  flame  —  is  unsatisfactory  because  so 
great  a  proportion  of  the  material  is  dispersed  before  the  heat  of 
the  flame  effects  the  dissociation  of  the  salt,  much  better  results 
may  be  obtained  by  making  use  of  more  powerful  flames  and 
substituting  for  the  single  loop  the  hollow  coils  of  platinum  wire 
first  recommended  by  Truchot  J  for  the  quantitative  determina- 
tion of  lithium.  Such  coils  are  easily  made  by  winding  the  wire 
somewhat  obliquely  about  a  rod  of  suitable  size,  pressing  the 

*  See  page  106. 

f  F.  A.  Gooch  and  T.  S.  Hart,  Am.  Jour.  Sci.,  [3],  xlii,  448. 

J  Compt.  rend.,  Ixxviii,  1022. 


THE  ALKALI   METALS 


81 


coils  close  together,  and  gathering  the  free  ends  into  a  twisted 
handle.  The  size  of  the  coils  is  adjustable  without  difficulty, 
so  that  each  coil  may  be  made  to  hold  almost  exactly  any  appro- 
priate amount,  and  to  take  up  this  amount  with  very  little  varia- 
tion in  successive  fillings,  provided  only  that  the  precaution  be 
taken  in  the  process  of  filling  to  plunge  the  coil  while  hot  into  the 
liquid,  and  to  keep  its  axis  inclined  obliquely  to  the  surface  of  the 
liquid  while  withdrawing  it. 

The  coils,  after  use,  may  be  conveniently  cleaned  by  heating 
them  in  the  flame  of  an  annular  burner  beneath  which  is  burned 
in  a  small  lamp  a  5  per  cent  solution  of  chloroform  in  alcohol, 
the  products  of  combustion  of  the  alcohol  and  chloroform  being 
conveyed  to  the  interior  of  the  flame  from  below  by  a  glass  funnel 
fitted  by  a  cork  to  the  tube  of  the  burner.  This  arrangement 
of  apparatus  gives  a  hot,  colorless  flame  through  which  hydro- 
chloric acid  is  constantly  diffused  in  condition  to  clean  the  wires 
completely  and  without  attention.  How  closely  the  capacity 
of  such  coils  may  be  adjusted,  and  how  uniformly  they  may  be 
filled,  is  shown  in  the  figures  of  the  accompanying  record. 

Capacity  of  Coils. 


I. 

grm. 

II. 

grm. 

III. 
grm. 

IV. 

grm. 

V. 
grm. 

VI.     . 

grm. 

Weight  of  filled  coil 

0.1996 
o.  1996 
0.1996 
0.1996 
0.1996 
0.1986 
O.OOIO 

0.2780 
0.2780 
0.2780 
0.2780 
0.2781 
o.  2760 

O.OO2O2 

0.2794 

0.2794 
0.2794 
0.2794 
0.2794 

o.  2764 
0.0030 

0.2844 
0.2845 
0.2844 
0.2845 
o.  2844 
o  .  2804 
o  .  00404 

0-3572 
0.3571 
0-3572 
0-3571 
0-3571 
0.3521 
o  .  00504 

0.3296 
0.3296 
0.3298 
0.3298 
0.3296 
0.3100 

o  01968 

Weight  of  filled  coil.  . 

Weight  of  filled  coil 

Weight  of  filled  coil  

Weight  of  filled  coil  

Weight  of  empty  coil  
Weight  of  contents  (mean).. 

It  is  plain  that  these  coils  afford  simple  means  of  taking  up 
known  amounts  of  material  in  solution.  By  gentle  heating  the 
liquid  may  be  evaporated  and  the  solid  material  left  thinly  spread 
and  in  condition  to  be  acted  upon  with  effect  when  brought  to 
the  flame.  The  evaporation  may  be  conducted  with  little  danger 
of  loss  of  material  by  holding  the  handle  of  the  coil  across  the 
flame  with  the  coil  proper  at  a  safe  distance  outside;  or,  prefer- 
ably, by  exposing  the  coils  over  a  hot  radiator,  the  handles  resting 
upon  a  flat  asbestos  ring. 

In  test  experiments  with  such  coils,  use  was  made  of  a  Muncke 
burner  giving  a  flame  3  cm.  wide  at  the  base  and  20  cm.  in  height. 


82  METHODS  IN  CHEMICAL  ANALYSIS 

The  coil  was  introduced,  after  thorough  drying,  just  within  the 
outer  mantle,  on  the  side  next  the  spectroscope,  with  the  axis 
transverse  to  the  slit  of  the  spectroscope  and  the  handle  across 
the  body  of  the  flame.  The  spectroscope  used  was  a  well-made 
single-prism  instrument  with  adjustable  slit  and  scale.  Observa- 
tions were  made  in  the  ordinary  diffused  light  of  the  laboratory, 
with  the  eye  in  use  shielded,  the  eye  not  in  use  covered,  and  the 
scale  illuminated  to  the  lowest  degree  of  visibility. 

Upon  experimenting  with  the  apparatus  described,  it  was  found 
that  coils  holding  0.02  grm.  of  water,  measuring  2  mm.  in  diam- 
eter by  i  cm.  in  length,  made  of  No.  28  wire  (0.32  mm.  in  diam- 
eter), and  wound  in  about  thirty  turns,  were  well  adapted  to  the 
purpose.  With  these  coils  and  the  flame  adjusted  to  a  height 
of  20  cm.,  7^  mg.  of  potassium  to  the  coilful  produces  a  line 
distinctly  visible  with  a  slit  of  0.18  mm.,  and  y^V^  mg-  with  a  slit 
of  0.23  mm.;  and  it  is  evident  that  this  practical  method  of 
producing  the  spectrum  of  potassium  gives  results  of  a  delicacy 
approaching  that  indicated  in  the  experiments  of  Bunsen  and 
Kirchhoff. 

These  determinations  were  made  with  pure  potassium  chlo- 
ride carefully  prepared  from  the  chlorate,  but  in  practical 
analysis  it  almost  always  happens  that  sodium  is  also  present. 
Experiments  were  therefore  made  to  determine  the  influence  of 
varying  amounts  of  the  latter  upon  the  visibility  of  the  potas- 
sium line.  The  dilution  of  the  potassium  chloride  was  adjusted 
nearly  to  the  last  limit  of  visibility,  so  that  a  coilful  of  the  liquid 
should  contain  7^  mg.,  or  yoW  mg-  °f  the  element,  according 
as  the  slit  was  0.18  mm.  or  0.23  mm.  wide;  to  this  solution  were 
added  weighed  amounts  of  pure  sodium  chloride  twice  repre- 
cipitated  and  washed  by  hydrochloric  acid;  and  the  spectro- 
scopic  tests  were  carried  out  as  before,  the  sodium  line  being 
kept  within  the  field  of  view  with  the  potassium  line. 

It  is  plain  from  the  results  (p.  83)  that  a  considerable  amount 
of  sodium  may  be  present  in  the  flame,  when  the  sodium  line 
is  in  full  view  in  the  spectrum,  and  the  slit  adjusted  to  nearly 
the  lowest  limit  of  visibility  of  pure  potassium,  without  inter- 
fering with  the  appearance  of  the  potassium  line,  but  that  a 
quantity  of  sodium  amounting  to  a  hundred  times  that  of  the 
potassium  may  be  sufficient  to  overpower  entirely  the  spectrum 
of  the  potassium.  The  inference  is  plain  that  the  proportion  of 


.THE  ALKALI  METALS 


sodium  to  potassium  should  not  be  permitted  to  reach  100:1 
when  it  is  desirable  to  bring  out  the  full  delicacy  of  the  spec- 
troscopic  test  with  the  sodium  line  in  the  field  of  view.  When 
too  great  a  proportion  of  sodium  is  present ,,  its  influence  may  be 
moderated  by  throwing  the  sodium  line  out  of  view,  if  the 
instrument  in  use  possesses  the  necessary  adjustment;  other- 
wise, it  is  easy  to  effect  a  partial  separation  of  the  sodium  chlor- 
ide from  the  potassium  chloride,  before  bringing  the  solution  to 
the  "test,  by  precipitating  with  alcohol,  and  experience  shows 
that  the  delicacy  of  the  test  for  potassium  is  not  impaired 
materially  by  such  treatment  of  the  mixed  chlorides. 

Effect  of  Sodium  upon  the  Potassium  Line, 


Weight  of  K 
in  a  coil- 
ful. 

mg. 

Weight  of 
Na  in  a 
coilful. 

mg. 

Ratio  of 
Na:K. 

Width  of 
slit. 

mm. 

Number  of 
trials. 

Characteristic  of  line. 

O.OOIO 

o.oooo 

o  : 

0.23 

3 

Visible. 

O.OOIO 

O.OO2O 

2 

0.23 

3 

Visible. 

O.OOIO 

O.OIOO 

10 

0.23 

3 

Visible. 

O.OOIO 

0.0200 

'    2O 

0.23 

3 

Visible. 

O.OOIO 

0.0400 

40 

0.23 

3 

Visible. 

O.OOIO 

O.O5OO 

SO 

0.23 

4 

Very  faint  or  none. 

O.OOIO 

O.  IOOO 

IOO 

0.23 

3 

None. 

O.OOIO 

O.2OOO 

2OO 

0.23 

3 

None. 

0.0014 

O.OOOO 

O 

0.18 

3 

Visible. 

0.0014 

0.0560 

40 

0.18 

3 

Visible. 

0.0014 

0.0700 

50 

0.18 

3 

Visible. 

0.0014 

0.1400 

IOO 

0.18 

2 

Visible. 

0.0014 

0.1400 

IOO 

0.18 

2 

None. 

Certain  experiments,  in  which  the  same  method  of  manipula- 
tion was  applied  to  the  determination  of  potassium  salts  other 
than  the  chloride,  indicate  that  the  test  is  less  delicate  in  the 
case  of  the  sulphate,  but  rather  more  delicate  in  the  case  of  the 
carbonate,  the  red  line  of  potassium  showing  unmistakably  in 
the  latter  case  when  only  ^Vs  of  a  milligram  of  potassium  was 
present. 

Determination  It  has  also  been  shown  by  Gooch  and  Hart  *  that 
of  Potassium,  ^e  quantitative  determination  of  small  amounts  of 
potassium  may  be  successfully  accomplished  by  bringing  an  un- 
known solution  of  potassium  chloride  to  the  measured  volume 

*  Loc.  cit. 


84  METHODS  IN  CHEMICAL  ANALYSIS 

at  which  the  residue  left  after  evaporating  the  contents  of  a  coil 
gives  a  line  of  the  same  strength  as  that  produced  by  the  residue 
of  a  coilful  of  a  standard  solution  of  potassium  chloride  and 
so  determining  the  concentration  of  the  unknown  solution  of 
finally  measured  volume.  It  is  convenient  to  use  several  coils 
adjusted  to  the  same  capacity,  and  to  clean,  fill,  dry  and  ignite 
them  before  the  spectroscope  in  the  manner  previously  de- 
scribed. From  time  to  time  the  capacity  of  the  coils  should 
be  readjusted,  or  else  final  comparison  tests  should  be  made 
with  a  single  coil.  It  is  essential  that  the  eye  of  the  observer 
should  be  kept  as  nearly  as  possible  in  the  same  condition  of 
sensitiveness  and  in  the  same  position  in  making  the  compari- 
sons, and  best  to  hold  the  eye  at  the  observing  telescope  during 
the  entire  interval  between  the  exposures,  shading  it  carefully, 
and  to  light  the  comparison  scale  of  the  spectroscope  to  the 
faintest  possible  visibility  sufficient  to  fix  exactly  the  position 
in  which  the  line  is  to  be  sought.  It  is  important,  too,  that  the 
trials  of  the  test  and  standard  should  come  as  closely  together 
as  possible  in  point  of  time.  The  observations  of  a  series  should 
be  made  by  the  same  individual,  the  preparation  and  exposure 
of  the  wires  being  made  by  another.  It  is  not  possible  to  attain 
the  best  results  in  such  work  single-handed.  The  dilution  of 
the  test  solution  is  made  conveniently,  and  with  sufficient 
accuracy,  in  ioo-cm.3  cylinders  graduated  to  half  cubic  centi- 
meters. It  is  advantageous  to  take  a  standard  solution  which 
corresponds  to  the  presence  of  yjg-  mg.  of  potassium  to  the  coil- 
ful, and  set  the  slit  at  a  width  sufficient  to  give  lines  for  com- 
parison bright  enough  to  be  visible  without  much  effort.  The 
mode  of  proceeding  is  to  dilute  the  test  solution  until  the  line 
given  by  the  potassium  contained  in  a  coilful  is  of  the  same 
brightness  as  that  given  by  the  same  quantity  of  the  standard 
solution.  From  the  final  volume  of  the  test-solution  the  quantity 
of  potassium  present  in  it  is  directly  calculable;  for,  since  any 
given  volume  of  the  test  solution  at  its  final  dilution  contains 
exactly  the  same  amount  of  potassium  as  the  same  volume  of  the 
standard  solution,  it  is  only  necessary  to  multiply  the  number 
expressing  the  volume  in  cubic  centimeters  of  the  test  solution 
by  that  of  the  weight  in  grams  of  the  potassium  contained  in  one 
cubic  centimeter  of  the  standard  in  order  to  obtain  the  weight  in 
grams  of  potassium  in  the  whole  test  solution. 


THE  ALKALI  METALS 


The  following  is  the  record  of  the  comparison  of  two  unknown 
test  solutions  of  pure  potassium  chloride  with  a  standard  solu- 
tion containing  o.oooi  grm.  of  the  same  salt  to  I  cm.3 

Determination  of  Potassium  in  Pure  Potassium  Chloride. 


Experiment  I. 

Experiment  II. 

Volume  of  test 

Characteristic  of  line 

Volume  of  test 

Characteristic  of  line 

solution. 

compared  with 

solution. 

compared  with 

standard. 

standard. 

cm.» 

cm.1 

2O 

Stronger. 

30 

Stronger. 

50 

Stronger. 

60 

Stronger. 

IOO 

Stronger. 

82 

Weaker. 

no 

Stronger. 

70 

Stronger. 

1  20 

Stronger. 

76 

Stronger. 

150 

Like. 

78 

Stronger. 

2OO 

Weaker. 

80 

Like. 

1  60 

Weaker. 

150 

Like. 

(150  X  o.oooi  =  0.0150) 


Potassium  found. 
Potassium  taken, 


Limits  on  either  side. . 
Error.. 


0.0150  grm. 
0.0150  grm. 
0.0120  grm. 
0.0160  grm. 
o.oooo  grm. 


(80  X  o .  oooi  =  o .  0080) 

Potassium  found 0.0080  grm. 

Potassium  taken 0.0080  grm. 

T  •    '^  -t.1  (  0.0078  grm. 

Limits  on  either  side    < 

I  0.0082  grm. 

Error o .  oooo  grm. 


These  results  with  the  pure  potassium  salt  show  a  degree  of 
accuracy  quite  unexpected.  In  the  former  no  attempt  was  made 
to  approximate  as  closely  as  possible  to  the  limits  of  dilution  on 
both  sides  of  the  condition  of  equal  brightness  in  test  and  stand- 
ard, but  in  the  latter  great  care  was  taken  in  this  respect  and  the 
.possible  error  does  not  exceed  two  and  a  half  per  cent  of  the  en- 
tire amount  of  potassium  involved.  Similar  experiments  with 
potassium  chloride  in  presence  of  varying  amounts  of  sodium 
chloride,  the  sodium  line  being  turned  from  the  field  of  view,  led 
to  a  recognition  of  the  fact  that  sodium  chloride  tends  to  increase 
the  brilliance  of  the  potassium  line,  —  the  maximum  strengthen- 
ing effect  of  about  20  per  cent  occurring  when  the  amount  of 
sodium  chloride  stands  to  that  of  the  potassium  in  the  ratio  of 
10  : 1,  —  a  phenomenon  due  to  the  chemical  action  of  sodium 
dissociated  in  the  flame.  The  effect  of  ammonium  chloride, 
and  of  hydrochloric  acid,  in  destroying  the  potassium  light  is 
well  known,  and  is  due,  presumably,  in  very  large  degree,  to 


86 


METHODS  IN  CHEMICAL  ANALYSIS 


prevention  of  the  dissociation  of  the  potassium  chloride.  The 
dissociated  sodium  should  naturally  reenforce  the  disintegrating 
action  of  heat  upon  the  potassium  chloride. 

The  complication  introduced  by  the  presence  of  any  certain 
amount  of  the  sodium  salt  in  the  test  may  be  obviated  by  the 
addition  of  the  same  amount  of  the  sodium  salt  to  the  standard, 
and  experience  shows  that  an  unknown  amount  of  the  sodium 
salt  in  the  test  may  be  matched  with  a  degree  of  accuracy  suffi- 
cient for  the  end  in  view.  The  determination  of  potassium  in 
the  presence  of  sodium  is  performed,  therefore,  in  three  stages: 
first,  the  test  solution  is  diluted  until  its  potassium  line  matches 
approximately  that  of  the  standard  made  to  contain  in  I  cm.3 
o.oooi  grm.  of  potassium  and  o.ooio  grm.  of  sodium  chloride; 
secondly,  sodium  chloride  is  added  to  the  solution  thus  diluted 
until  the  sodium  lines  of  test  and  standard  are  brought  to  equal- 
ity;  and,  finally,  the  potassium  lines  of  test  solution  and  standard 
solution  are  again  brought  into  comparison.  Following  are  the 
records  of  experiments  made  in  this  manner. 


Determination  of  Potassium  in  Presence  of  Sodium. 
Experiment  I. 


Parti. 

Part  II. 

Part  III. 

Volume 
of  test 
solu- 
tion 

Width 
of  slit. 

Character- 
istic of 
potassium 
line  as  com- 

NaCl in 
100  cm.8 
of  test 
solu- 

Width 
of  slit. 

Character- 
istic of 
sodium  line 
as  compared 

Volume 
of  test 
solu- 

Width 
of  slit. 

Character- 
istic of 
potassium 
line  as  com- 

pared with 

tion. 

with 

pared  with 

cm.* 

mm. 

standard. 

grm. 

mm. 

standard. 

cm.3 

mm. 

standard. 

30 
70 
100 

0.23 
0.23 
0.23 

Stronger. 
Stronger. 
Weaker. 

0.01* 
0.03 
O.O5 

OO  GO  OO 
M  M  M 

6  d  d 

Weaker. 
Weaker. 
Weaker. 

108 
108 

0.23 
0.23 

'Weaker. 
Stronger. 
(Weaker. 

0.08 
0.09 

0.18 
0.18 

Weaker. 
Weaker. 

109 

0.23 

}  Stronger. 
/Like. 

O.IO 

0.18 

Like. 

*  Originally  present.  v 

The  test  solution  having  been  accidently  over-diluted,  its  strength  was 
increased  by  the  addition  of  o.ooio  grm.  of  potassium,  and  this  amount  was 
added  in  the  computation  below  to  that  originally  in  the  test  solution. 
(109  X  o.oooi  =  0.0109) 

Potassium  found 0.0109  grm. 

Potassium  taken. o.ono  grm. 

Error o .  oooi  grm.  =0.9  per  cent. 


THE  ALKALI  METALS 


Determination  of  Potassium  in  Presence  of  Sodium. 

Experiment  II. 


Part  I. 

Part  II. 

Part  III. 

Volume 
of  test 
solu- 

Width 
of  slit. 

Character- 
istic of 
potassium 
line  as  com- 

NaCl in 
100  cm.3 
of  test 
solu- 

Width 
of  slit. 

Character- 
istic of 
sodium  line 
as  compared 

Volume 
of  test 
solu- 
tion. 

Width 
of  slit. 

Character- 
istic of 
potassium 
line  as  com- 

pared with 

tion. 

with 

pared  with 

cm.3 

mm. 

standard. 

grm. 

mm. 

standard. 

cm.8 

mm. 

standard. 

40 
100 

0.23 
0.23 

Stronger. 
Stronger. 

0.025* 
0.050 

0.18 
0.18 

Weaker. 
Weaker. 

1  60 
1  80 

0.23 
0.23 

Stronger. 
Stronger. 

160 

0.23 

Weaker. 

0.085 

0.18 

Weaker. 

I  QO 

0.23 

Stronger. 

O.IOO 

0.18 

Weaker. 

2OO 

0.23 

Stronger. 

O.IIO 

0.18 

Like. 

205 

0.23 

Weaker. 

2IO 

0.23 

Weaker. 

*  Originally  present. 

/  205  X  o.oooi  =0.0205  )  \ 

>  mean  =  0.02021;  1 

\2OO  X  O.OOOI   =O.O2OO  )  / 

Potassium  found 0.02025  grm- 

Potassium  taken 0.02000  grm. 

Error 0.00025  grm.  =  1.25  per  cent. 


Experiment  III. 


Parti. 

Part  II. 

Part  III. 

Volume 
of  test 
solu- 
tion. 

Width 
of  slit. 

Character- 
istic of 
potassium 
line  as  com- 

NaCl in 
ico  cm.3 
of  test 
solu- 

Width 
of  slit. 

Character- 
istic of 
sodium  line 
as  compared 

Volume 
of  test 
solu- 
tion. 

Width 
of  slit. 

Character- 
istic of 
potassium 
line  as  com- 

cm.3 

mm. 

pared  with 
standard. 

tion, 
grm. 

mm. 

with 
standard. 

cm.3 

mm. 

pared  with 
standard. 

40 
80 
IOO 

0.23 
0.23 
0.23 

Stronger. 
Stronger. 
Stronger. 

0.045* 
0.082 

0.18 
0.18 

Weaker. 
Like. 

no 

120 
130 

6.23 
0.23 
0.23 

Stronger. 
Stronger. 
Like. 

no 

0.23 

Like. 

*  Originally  present. 

(130  X  o.oooi  =  0.0130) 

Potassium  found 0.0130  grm. 

Potassium  taken 0.0140  grm. 

Error o.ooio  grm.  =  7  per  cent. 


88 


METHODS  IN  CHEMICAL  ANALYSIS 


Determination  of  Potassium  in  Presence  of  Sodium. 

Experiment  IV. 


Part  I. 

Part  II. 

Part  III. 

Volume 
of  test 
solu- 
tion. 

Width 
of  slit. 

Character- 
istic of 
potassium 
line  as  com- 
pared with 

NaCl  in 
100  cm.3 
of  test 
solu- 
tion. 

Width 
of  slit. 

Character- 
istic of 
sodium  line 
as  compared 
with 

Volume 
of  test 
solu- 
tion. 

Width 
of  slit. 

Character- 
istic of 
potassium 
line  as  com- 
pared with 

cm.* 

mm. 

standard. 

grm. 

mm. 

standard. 

cm.* 

mm. 

standard. 

First. 

7rt 

022 

Stronger 

o  18 

Weaker 

90 

100 

0.23 
0.23 

Stronger. 
(  Weaker. 
\Like. 

0.07 
0.09 

O.IO 

0.80 
0.18 

0.18 

Weaker. 
Weaker. 
(Like. 
<  Stronger. 
(  Stronger. 

100 
1  2O 
130 

140 

0.23 
0.23 
0.23 

0.23 

Stronger. 
Stronger. 
Stronger. 
(  Stronger. 
I  Weaker. 

Second. 

120 

0.23 

Stronger. 

140 

0.23 

Stronger. 

j  Stronger. 

•% 

I5° 

0.23 

{  Weaker. 

1  60 

0.23 

Weaker. 

*  Originally  present. 

First.  Second. 

(140  X  o.oooi  =  0.0140)  (150  X  o.oooi  =  0.0150) 

Potassium  found ...  0.0180  grm.  0.0150 

Potassium  taken.  . .  0.0150  grm.  0.0150 

Error o.ooio  grm.  =  7  per  cent o.oooo 

Though  not  accurate  in  the  highest  degree  when  considerable 
amounts  of  potassium  are  to  be  estimated,  the  method  is  reason- 
ably applicable  to  the  determination  of  small  quantities  of  that 
element. 


The  Separation  and  Determination  of  Potassium  as  the  Perchlorate. 

Kreider's  method*  for  the  preparation  of  perchloric  acid  has 
greatly  facilitated  the  use  of  the  perchlorate  method  for  the 
estimation  of  potassium.  This  method  —  consisting  essentially 
in  the  regulated  heating  of  sodium  chlorate  (readily  obtained 
in  the  market) ,  treatment  of  the  residue  with  strong  hydrochloric 
acid  to  yield  a  precipitate  of  sodium  chloride  and  a  solution  of 
perchloric  acid  containing  a  small  amount  of  sodium  chloride, 
*  D.  Albert  Kreider,  Am.  Jour.  Sci.,  [3],  xlix,  443. 


THE   ALKALI  METALS  89 

and  filtration  of  the  mixture  upon  asbestos  —  may  be  detailed 
as  follows : 

A  convenient  quantity  of  sodium  chlorate,  from  100  to  300 
grm.,  is  melted  in  a  glass  retort  or  round-bottomed  flask  and 
gradually  raised  to  a  temperature  at  which  oxygen  is  freely  but 
not  too  rapidly  evolved,  and  kept  at  this  temperature  for  one 
and  a  half  or  two  hours,  until  the  thickening  of  the  mass  indi- 
cates  the  conversion  of  the  chlorate  to  chloride  and  perchlorate; 
or,  the  retort  may  be  connected  with  a  gasometer  and  the  end 
of  the  reaction  determined  by  the  volume  of  oxygen  expelled, 
according  to  the  equation 

2  NaClO3  =  NaCl  +  NaClO4  +  O2. 

The  product  thus  obtained  is  washed  from  the  retort  to  a  capa- 
cious evaporating  dish,  where  it  is  treated  with  sufficient  hydro- 
chloric acid  to  effect  the  complete  reduction  of  the  residual 
chlorate,  which,  if  the  ignition  has  been  carefully  conducted 
with  well-distributed  heat,  will  be  present  in  but  small  amount. 
It  is  then  evaporated  to  dry  ness  on  the  steam  bath,  or  more 
quickly  over  a  direct  flame,  and  with  but  little  attention  until  a 
point  near  to  dry  ness  has  been  reached.  Then  stirring  is  found 
of  great  advantage  in  facilitating  the  volatilization  of  the  remain- 
ing liquid  and  in  breaking  up  the  mass  of  salt;  otherwise  the 
perchlorate  seems  to  solidify  with  a  certain  amount  of  water, 
and  removal  from  the  dish,  without  moistening  and  reheating, 
is  impossible. 

After  trituration  of  the  residue  in  a  porcelain  mortar,  an  ex- 
cess of  the  strongest  hydrochloric  acid  is  added  to  the  dry  salt, 
preferably  in  a  tall  beaker,  where  there  is  less  surface  for  the 
escape  of  hydrochloric  acid  and  from  which  the  acid  may  be 
decanted  without  disturbing  the  precipitated  chloride.  If  the 
salt  has  been  reduced  to  a  very  fine  powder,  by  stirring  ener- 
getically for  a  minute,  the  hydrochloric  acid  will  set  free  the 
perchloric  acid  and  precipitate  the  sodium  as  chloride,  which  in 
a  few  minutes  settles,  leaving  a  clear  solution  of  the  perchloric 
acid  with  the  excess  of  hydrochloric  acid.  The  clear  supernatant 
liquid  is  then  decanted  upon  asbestos  in  a  perforated  crucible, 
through  which  it  may  be  rapidly  drawn  with  the  aid  of  suction, 
and  the  residue  is  again  treated  with  the  strongest  hydrochloric 
acid.  The  liquid  is  again  decanted,  the  salt  is  finally  brought 


90  METHODS  IN  CHEMICAL  ANALYSIS 

upon  the  filter,  where  it  is  washed  with  a  little  strong  hydro- 
chloric acid.  A  large  platinum  cone  will  be  found  more  con- 
venient than  the  crucible,  because  of  its  greater  capacity  and 
filtering  surface. 

The  filtrate,  containing  the  perchloric  acid  with  the  excess  of 
hydrochloric  acid  and  the  small  per  cent  of  sodium  chloride  which 
is  soluble  in  the  latter,  is  then  evaporated  over  the  steam  bath 
till  all  hydrochloric  acid  is  expelled  and  the  heavy  white  fumes 
of  perchloric  acid  appear,  when  it  is  ready  for  use  in  potassium 
determinations.  Evidently  the  acid  is  not  chemically  pure  be- 
cause the  sodium  chloride  is  not  absolutely  insoluble  in  hydro- 
chloric acid ;  but  a  test  with  silver  nitrate  proves  that  the  sodium, 
together  with  any  other  bases  which  may  have  gone  through  the 
filter,  has  been  completely  converted  into  perchlorate,  and,  unless 
the  original  chlorate  contained  potassium  or  the  acid  had  been 
exposed  to  the  fumes  of  ammonia,  the  residue  of  evaporation, 
which  does  not  exceed  0.04  grm.  in  weight  to  I  cm.,3  is  easily  and 
completely  soluble  in  97  per  cent  alcohol.  Perchloric  acid  thus 
prepared  was  found  to  contain  0.9831  grm.  of  free  anhydrous 
acid  in  I  cm.3. 

Should  the  sodium  chlorate  used  in  the  process  contain  potas- 
sium as  an  impurity,  the  mixture  of  sodium  perchlorate  and 
chloride,  after  being  treated  with  hydrochloric  acid  for  the  reduc- 
tion of  the  residual  chlorate,  is  reduced  to  a  fine  powder,  and  well 
digested  with  97  per  cent  alcohol,  which  dissolves  the  sodium 
perchlorate,  but  leaves  the  chloride  as  well  as  any  potassium  salt 
insoluble.  The  alcoholic  solution  of  the  perchlorate  is  then  dis- 
tilled from  a  large  flask  until  the  perchlorate  begins  to  crystallize, 
when  the  heat  is  removed  and  the  contents  quickly  emptied  into 
an  evaporating  dish.  The  mixture  is  evaporated  to  dryness  on 
the  steam  bath  and  the  residue  is  treated  with  strong  hydro- 
chloric acid  for  the  separation  of  the  perchloric  acid  in  the  manner 
described  above. 

In  applying  perchloric  acid,  prepared  by  Kreider's  method, 
to  the  determination  of  potassium  according  to  the  treatment 
suggested  by  Caspari,*  very  satisfactory  results  were  obtained. 
Briefly,  the  method  is  as  follows;  The  substance,  free  from  sul- 
phuric acid,  is  evaporated  to  the  expulsion  of  free  hydrochloric 
acid.  The  residue,  stirred  with  20  cm.3  of  hot  water  and  then 
*  Zeit.  angew.  Chem.,  1893,  68. 


THE  ALKALI  METALS 


treated  with  perchloric  acid  in  quantity  not  less  than  one  and 
one-half  times  that  required  by  the  bases  present,  is  evaporated 
with  frequent  stirring  to  a  thick,  sirup-like  consistency,  again  dis- 
solved in  hot  water  and  evaporated  with  continued  stirring  until 
all  hydrochloric  acid  has  been  expelled  and  the  fumes  of  per- 
chloric acid  appear.  Further  loss  of  perchloric  acid  is  to  be  com- 
pensated for  by  addition  of  more.  The  cold  mass  is  then  well 
stirred  with  about  20  cm.3  of  wash  alcohol  —  97  per  cent  alcohol 
containing  0.2  per  cent  by  weight  of  pure  perchloric  acid  —  with 
precautions  against  reducing  the  potassium  perchlorate  crystals 
to  too  fine  a  powder.  After  settling,  the  alcohol  is  decanted  on 
the  asbestos  filter  and  the  residue  is  again  washed  with  about 
the  same  amount  of  wash  alcohol.  The  residual  salt,  freed 
from  alcohol  by  gently  heating,  is  dissolved  in  10  cm.3  of  hot 
water  and  a  little  perchloric  acid.  The  solution  is  evaporated 
once  more  with  stirring  until  fumes  of  perchloric  acid  rise.  The 
precipitate  is  treated  with  I  cm.3  of  wash  alcohol,  transferred  to 
the  asbestos,  preferably  by  a  policeman,  and  washed  with  pure 
alcohol ;  the  whole  process  requiring  about  50  to  70  cm.3  of  alco- 
hol. It  is  then  dried  at  about  130°  C.  and  weighed. 

The  substitution  of  a  perforated  crucible  for  the  truncated 
pipette  employed  by  Caspari  is  advantageous;  and  asbestos 
capable  of  forming  a  close,  compact  felt  should  be  selected , 
inasmuch  as  the  perchlorate  is  in  part  unavoidably  reduced, 
during  the  necessary  stirring,  to  so  fine  a  condition  that  it  tends 
to  run  through  the  filter  when  under  pressure. 

A  number  of  determinations  made  of  potassium  unmixed  with 
other  bases  or  nonvolatile  acids  are  recorded  in  the  following 

table : 

Potassium  in  the  Pure  Salt. 


KC1  taken. 

Volume  of 
filtrate. 

KC1O4  found. 

Error  on, 
KC1O4. 

Error  on  KC1. 

Error  on  K,O. 

grm. 

cm.3 

grm. 

grm. 

grm. 

grm. 

0.  IOOO 

54 

0.1851 

—0.0008 

—  0.0004 

—  0.0003 

O.  IOOO 

58 

0.1854 

—0.0005 

—  O.OOO2 

—  O.OOO2 

O.  IOOO 

5i 

0.1859 

o.oooo 

O  .  OOOO 

O  .  OOOO 

0.1000 

50 

0.1854 

—0.0005 

—  O.OOO2 

—  O.OOO2 

O.IOOO 

48 

0.1859 

0.0000 

0.0000 

O.OOOO 

O.  IOOO 

52 

0.1854 

—  0.0005 

—  0.0002 

—  O  .  OOO2 

As  Caspari  has  pointed  out,  sulphuric  acid  must  be  removed 
by  precipitation  as  barium  sulphate  before  the  treatment  with 


92 


METHODS  IN   CHEMICAL  ANALYSIS 


perchloric  acid  is  attempted,  and  unless  the  precipitation  is  made 
in  a  strongly  acid  solution,  some  potassium  is  carried  down  with 
the  barium.  Phosphoric  acid  need  not  be  previously  removed; 
but  to  secure  a  nearly  complete  separation  of  this  acid  from  the 
potassium,  a  considerable  excess  of  perchloric  acid  should  be 
left  upon  the  potassium  perchlorate  before  it  is  treated  with  the 
alcohol.  When  these  conditions  are  carefully  complied  with, 
fairly  good  results  may  justly  be  expected.  Below  are  given  a 
number  of  the  results  obtained. 

Potassium  in  Mixtures  of  Salts. 


Compounds  taken, 
grin. 

Volume 
of 
filtrate. 

cm.8 

KC104 
found. 

grm. 

Error  on 
KC1O4. 

grm. 

Error  on 
KC1. 

grm. 

Error  on 
K20. 

grm. 

KCl  -o.ioool 

CaCO3=o.i3 
MgSO4  =  o.i3 
Fe2Cl6  =0.05      ^ 
Al2{S04)3  =  o.o5 
MnO2=o.o5      | 
HNa2PO4.i2H20  =  o.4o     J 

50 
82 
80 
80 
92 
60 

0.1887 
0.1875 
0.1861 
0.1843 
0.1839 
0.1854 

+0.0028 
+O.OOl6 
+0.0002 

—  0.0016 

—  0.0020 
—  O.OOO5 

+0.0014 
+0.0008 

+O.OOOI 

—0.0008 

—  O.OOIO 
—  O.OOO2 

+0.0009* 
+0.0005* 
+0.000lf 

—  0.0005! 
—  0.0006  t 

—  O.OOO2f 

*  The  residue  showed  phosphoric  acid  plainly  when  tested, 
t  Only  traces  of  phosphoric  acid  found  in  the  residue. 

In  the  last  three  experiments,  in  which  the  amount  of  perchloric 
acid  employed  was  about  three  times  that  required  to  unite 
with  the  bases  present,  the  phosphoric  acid  subsequently  found 
with  the  potassium  was  hardly  enough  to  appreciably  affect  the 
weight. 

The  Estimation  of  Potassium  as  the  Pyrosulphate. 

Browning*  has  shown  that  potassium  sulphate,  like  sodium 
sulphate  f  when  treated  with  sulphuric  acid  and  submitted  to 
a  temperature  ranging  between  250°  and  270°,  takes  very  defi- 
nitely the  form  of  the  pyrosulphate,  K2S2O7,  and  that  potassium 
may  be  estimated  as  that  salt.  Under  similar  conditions  caesium 
and  rubidium  remain  in  the  form  of  acid  sulphates.  Results  of 
the  procedure  are  given  below  in  comparison  with  the  results 
obtained  upon  moistening  the  pyrosulphate  with  ammonium 
hydroxide  and  igniting  strongly  to  form  the  neutral  sulphate. 

*  Philip  E.  Browning,  Am.  Jour.  Sci.,  [4],  xii,  301. 
t  See  page  79. 


THE  ALKALI  METALS 


93 


Potassium  Sulphates  by  Ignition. 


KCl  taken. 

grin. 

K,S20, 
calculated. 

grm. 

K2S207 
found. 

grm. 

Error, 
grm. 

K2S04 

calculated. 

grm. 

K2S04 
found. 

grm. 

Error, 
grm. 

O    2172 

O    37O4. 

o  3608 

—  o  0006 

o  .  i  706 

O.II92 
O.IO74 

o  1006 

0.2909 
0.2032 
0.1830 

o  1868 

0.2886 
O.2O22 
0.1823 

•o  1860 

-0.0023 
—  O.OOIO 

—  0.0007 
—  o  0008 

0.1993 
0-1393 

0.1972 
0.1381 

—  O.OO2I 
—  O.OOI2 

The  Volumetric  Estimation  of  Potassium  as  the  Cobalti-nitrite. 

The  use  of  sodium  cobalti-nitrite  to  estimate  potassium  has 
been  described  by  R.  H.  Adie  and  T.  B.  Wood,*  who  show  results 
fairly  accurate  and  favorably  comparable  with  those  obtained 
by  the  platinic  chloride  gravimetric  method.  In  the  process 
worked  out  by  these  investigators  a  solution  of  a  potassium  salt 
containing  the  equivalent  of  0.5  per  cent  to  I  per  cent  of  K2O 
is  acidified  with  acetic  acid  and  precipitated  by  an  excess  of 
sodium  cobalti-nitrite.t  The  mixture  is  allowed  to  stand  at  least 
a  few  hours,  preferably  over  night,  and  is  then  filtered  through  a 
perforated  crucible  fitted  with  an  asbestos  felt.  The  precipitate 
is  washed  with  10  per  ctent  acetic  acid.  According  to  Sutton, 
it  is  important  that  the  precipitation  should  be  made  in  a  solu- 
tion containing  the  equivalent  of  0.5  per  cent  to  I  per  cent  of 
K2O,  since  in  solutions  of  lower  concentration  the  precipitate 
conies  down  in  a  condition  in  which  it  is  apt  to  run  through  the 
filter  in  washing.  The  precipitate  is  then  decomposed  by  boiling 
in  dilute  sodium  hydroxide,  and  the  cobalt  is  removed  as  the 
hydroxide  by  filtration.  The  nitrites,  which  are  a  measure  of 
the  potassium  in  the  precipitate,  are  estimated  by  titrating  with 
standard  potassium  permanganate.  Adie  and  Wood  found  by 
analysis  that  the  composition  of  the  potassium  salt  precipitated 
in  presence  of  the  excess  of  sodium  cobalti-nitrite  is  represented 
by  the  formula  K2NaCo(NO2)6.H2O,  and  that  in  their  method 
a  cubic  centimeter  of  strictly  n/io  potassium  permanganate  is 
equivalent  to  0.000785  grm.  of  K2O. 

This  process  has  been  studied  by  DrushelJ  with  a  view  to 
determining  the  best  conditions  for  precipitating  and  filtering 

*  Jour.  Chem.  Soc.,  Ixxvii,  1076;  Button's  Vol.  Anal.,  9th  ed.,  page  62. 

f  Ibid. 

J  W.  A.  Drushel.  Am.  Jour.  Sci..  [4],  xxfv,  433. 


94  METHODS  IN  CHEMICAL  ANALYSIS 

the  potassium  cobalti-nitrite,  and  to  shortening  the  work  by 
oxidizing  the  precipitated  cobalti-nitrite  with  potassium  per- 
manganate without  the  preliminary  decomposition  of  the  precipi- 
tate and  removal  of  cobalt  recommended  by  Adie  and  Wood, 
the  excess  of  permanganate  being  reduced  by  standard  oxalic 
acid,  and  the  remaining  oxalic  acid  titrated  to  color.  In  this 
treatment  trivalent  cobalt  is  reduced  to  the  bivalent  condition; 
the  oxygen  thus  made  available  is  equivalent  to  one-twelfth  of 
that  necessary  to  oxidize  the  nitrites.  The  factor  used,  there- 
fore, in  calculating  the  results  from  the  direct  titration  should 
be  twelve-elevenths  of  that  given  by  Adie  and  Wood;  that  is, 
in  titrating  the  precipitate  without  first  separating  the  cobalt 
one  cubic  centimeter  of  strictly  n/io  potassium  permanganate 
is  equivalent  to  0.000857  grm-  °f  K2O. 

By  repeated  experiments  it  was  found  that  difficulty  in  filtra- 
tion, as  well  as  the  necessity  for  allowing  the  precipitate  to  stand 
over  night,  may  be  avoided  by  evaporating  the  mixture  nearly 
to  dryness  on  the  steam  bath  after  adding  the  sodium  cobalti- 
nitrite  solution  in  considerable  excess.  Upon  cooling  the  pasty 
residue  becomes  hard  and  dry.  When  treated  with  cold  water 
the  excess  of  sodium  cobalti-nitrite  dissolves,  and  the  insoluble 
portion  may  be  collected  and  freely  washed  without  showing  a 
tendency  to  pass  through  the  filter. 

Potassium  in  the  The  application  of  the  cobalti-nitrite  method  as 
Pure  Salt.  worked  out  by  Drushel  is  as  follows:  The  solution 
of  a  potassium  salt,  containing  not  more  than  0.2  grm.  K2O  and 
free  from  ammonium  salt,  is  treated  with  a  rather  large  excess 
of  sodium  cobalti-nitrite  solution  acidified  with  acetic  acid,  and 
evaporated  to  a  pasty  condition  over  the  steam  bath.  It  is 
then  cooled,  treated  with  50  cm.3  to  100  cm.3  of  cold  water  and 
stirred  until  the  excess  of  sodium  cobalti-nitrite  is  dissolved, 
allowed  to  settle,  and  decanted  through  a  perforated  crucible 
fitted  with  an  asbestos  felt.  The  precipitate  is  washed  two  or 
three  times  by  decantation,  after  which  it  is  transferred  to  the 
crucible  and  thoroughly  washed  with  cold  water.*  In  the  mean- 
time a  measured  excess  of  standard  potassium  permanganate  is 
diluted  to  ten  times  its  volume  and  heated  nearly  to  boiling. 

*  It  was  found  later  that  a  half-saturated  sodium  chloride  solution  is  pref- 
erable to  cold  water  for  washing  the  precipitate,  since  it  permits  the  use  of  a 
coarser  asbestos  felt  in  filtering  without  danger  of  loss, 


THE  ALKALI  METALS 


95 


Into  this  the  precipitate  and  felt  are  transferred  and  stirred,  after 
which  the  crucible  is  also  put  into  the  solution,  since  particles 
of  the  precipitate  stick  persistently  to  the  sides  of  the  crucible. 
After  the  oxidation  has  proceeded  five  or  six  minutes  manganese 
hydroxide  separates  out  and  the  color  of  the  solution  darkens. 
At  this  point  5  cm.3  to  25  cm.3  of  sulphuric  acid  [1:7]  are  added, 
and  the  solution,  after  stirring,  is  allowed  to  stand  a  few  minutes. 
Then  a  measured  amount  of  standard  oxalic  acid,  containing 
50  cm.3  of  strong  sulphuric  acid  per  liter,  is  run  in  from  a  burette, 
with  care  to  add  an  excess.  The  temperature  is  maintained 
a  little  below  the  boiling  point  until  the  solution  becomes  color- 
less and  the  manganese  hydroxide  has  completely  dissolved. 
Titration  is  then  effected  by  permanganate  in  the  usual  manner. 
The  results  of  the  experimental  tests  with  potassium  chloride 
alone  and  in  presence  of  salts. of  the  calcium  group  are  given 

below. 

Potassium  in  Pure  Potassium  Chloride. 


KoO  taken  as 
KC1. 

grtn. 

K2O  found. 

Error  in  K2O. 

Gravimetrically. 
grm. 

Volumetrically. 

grm. 

Gravimetrically  . 
grm. 

Volumetrically. 
grm. 

0.0237 
0.0237 

0-0354 
0.0474 
0.0048 
0.0024 
0.0005 
0.0015 
0.0355 

0.0240 
0.0243 

0-0359 
0.0478 
0.0048 
0.0024 

0.0238 
0.0242 
0-0355 
0.0471 
0.0050 
0.0023 
o  .  0006 

O.OOI7 
0-0355 

+0  .  0003 
+O.OOO6 
+O.OOO4 
+0.0004 
0.0000 

o.oooo 

+O.OOOI 

+o  .  0005 
o.oooo 
—0.0003 

+O.OOO2 
+O.OOOI 
+0.0001 

+0.0002 
o.oooo 

In  the  first  six  experiments  of  this  series  the  precipitate  was 
dried  at  115°,  weighed,  and  then  treated  with  permanganate. 

Potassium  in  Mixtures  of  Salts. 


CaCl2. 
grm. 

3S: 

grm. 

BaCl2  . 
taken. 

grm. 

Sr(N03)2. 
grm. 

K2O  taken, 
grm. 

K2O  found, 
grm. 

Error, 
grm. 

O    2OOO 

O    2OOO 

o  0005 

o  0007 

+O   OOO2 

o  3000 

o  5000 

o  0237 

o  0234 

—  o  0003 

o  5000 

I    OOOO 

o  0829 

o  0824 

—  o  0005 

o  5000 

I  .OOOO 

0.5000 

O.O7II 

o  0737 

+o  0026 

0.5000 

o  .  5000 

I  .OOOO 
I  .OOOO 

0.5000 
0.5000 

0.5000 

0.0474 
0.0237 

0.0493 

0.0251 

+O.OOI9 
+O.OOI4 

0.5000 

I  .OOOO 



0.07II 

0.0713 

+O.OOO2 

96 


METHODS  IN  CHEMICAL  ANALYSIS 


The  salts  of  calcium  and  magnesium  4o  not  influence  the 
accuracy  of  the  process,  while  the  presence  of  salts  of  barium 
and  strontium  tends  to  high  results. 

Potassium  in  The  method*  is  applicable  to  the  estimation  of 
Fertilizers.  potassium  in  fertilizers.  The  process,  as  laid  down, 
is  as  follows : 

Ten  grams  of  the  fertilizer  are  placed  in  a  5<x>-cm.3  flask, 
300  cm.3  of  water  added,  the  contents  boiled  for  30  minutes, 
and  ammonia  water  added  to  slight  alkalinity.  Enough  am- 
monium oxalate  is  added  to  precipitate  all  the  calcium,  and, 
after  cooling,  the  solution  is  made  up  to  the  mark  on  the  neck 
of  the  flask  and  well  shaken.  The  solution  is  then  filtered 
through  a  dry  filter  into  a  dry  flask,  and  5O-cm.3  portions  of 
the  filtrate  are  transferred  with  a  pipette  to  platinum  dishes,  for 
estimation  by  the  cobalti-nitrite  method.  After  evaporating 
these  portions  to  half  their  volume  over  the  steam  bath,  I  cm.3  of 
sulphuric  acid  [i:  i]  is  added  and  the  evaporation  is  continued 
as  far  as  possible  over  the  steam  bath,  and  finally  over  a  low 
flame.  After  the  danger  of  spattering  is  over,  the  flame  is  in- 
creased and  the  charred  organic  matter  is  burned  off,  finally, 
over  the  blast  lamp.  The  potassium  sulphate  is  dissolved  by 
adding  a  little  water  and  heating  over  the  steam  bath,  and  the 
potassium  is  estimated  as  in  the  previously  described  treatment 
of  the  pure  potassium  salt.f 

Potassium  in  Mixed  Fertilizers. 


KZO  by  platinum  chloride  method. 

K2O  by  cobalti- 

Water-soluble 

Number. 

nitrite  method. 

PzO5  in  sample. 

Per  cent. 

Per  cent. 

Per  cent. 

Per  cent. 

I 

5-22 

S-l8 

5-i8 

4.16 

2 

6-53 

6.56 

6.56 

3.10 

3 

2.23 

2.24 

2.24 

7.82 

4 

8.68 

8.64 

8.78 

o  94 

5 

6.37 

6.42 

'        6.38 

6.62 

6 

6.08 

6.13 

6.13 

S-6i 

7 

4.08 

4.02 

4.02 

3-iS 

8 

4.62 

4.66 

4.67 

2-43 

9 

1.68 

1.67 

1.77 

6.03 

*  W.  A.  Drushel,  Am.  Jour.  Sci.,  [4],  xxiv,  437. 
t  See  page  94. 


THE  ALKALI  METALS 


97 


In  the  preceding  table  are  given  results  obtained  by  the  method 
with  nine  fertilizers,  and,  for  comparison,  results  (by  two  ana- 
lysts) by  the  platinic  chloride  method. 

Potassium  in          In  applying  the  cobalti-nitrite  method  to  the  esti- 
Soiis.  mation  of  potassium  in  soils,  the  general  procedure 

may  be  outlined  as  follows :  * 

A  weighed  amount  of  dry  soil  is  extracted  with  an  excess  of 
hydrochloric  acid  over  the  steam  bath.  The  excess  of  acid  is 
removed  from  the  extract  by  evaporation.  The  bases  which 
might  interfere  with  the  process  are  removed  with  sodium  car- 
bonate or  ammonium  hydroxide  and  ammonium  oxalate.  Am- 
monium salts  and  organic  matter  are  removed  by  ignition. 

Potassium  in  Soils. 


Character  of  soil. 

Soil  taken, 
grm. 

K2O  found. 

Platinum 
chloride 
method. 

grm. 

Cobalti-nitrite 
method. 

grm. 

Per  cent. 

Clay  
Cfay  
Loam  

Loam  
Gravel  
Gravel.  .  .  . 

Clay 
gravel  .  . 

(d)  

](2)  

I  (0 

2-5 

2.5 
2.5 
2.5 
2.5 
2.5 
2.5 

2-5 

2.5 
2.5 
2.5 
2.5 
2.5 
2.5 
2.5 
2.5 
2.5 
2.5 
2.5 
2.5 
2.5 
2.5 

0.0028 
0.0035 

O.  II     , 

0.14 
0.14 
0-39 
0-37 
0-37 
0.30 
0.27 
0.30 
0.24 
0.23 
0.23 
0.17 
0.18 
o.  19 
0.18 

0.20 

0.19 
0.18 
0.18 
0.16 
0.18 

0.0035 
0.0093 
0.0075 
0.0058 

'((I).... 

1(3).. 

O.OIOO 

0.0092 

d)...  " 

0.0074 
0.0068 

v   (s) 

(d)    ... 

o  .  0060 
0.0058 

(  (i).. 

o  .  0042 

1(2).. 

0.0045 

(    (i)  

0.0047 
0.0044 

0.0050 
0.0046 

{d).. 

0.0048 

Si.; 

(3).. 

0.0045 
o  .  0040 

0.0044 

(4).. 

(s) 

Small  amounts  of  phosphoric  acid  do  not  interfere.     The  resi- 
due is  dissolved  in  a  little  water  and  a  few  drops  of  acetic  acid, 

*  W.  A.  Drushel,  Am.  Jour.  Sci.,  [4],  xxvi,  329. 


98  METHODS  IN  CHEMICAL  ANALYSIS 

and  the  mixture  evaporated  with  an  excess  of  sodium  cobalti- 
ni trite  to  a  pasty  condition,  stirred  up  with  cold  water,  and 
filtered  upon  asbestos  in  a  perforated  crucible.  The  precipitated 
potassium  sodium  cobalti-nitrite  is  washed  with  a  half-satu- 
rated solution  of  sodium  chloride,  and  treated  with  an  excess 
of  permanganate  in  hot  dilute  solution.  The  color  of  the  per- 
manganate is  destroyed  by  an  excess  of  standard  acidulated 
oxalic  acid,  and  the  excess  of  oxalic  acid  titrated  to  color  with 
permanganate. 

The  test  analyses  show  an  excellent  agreement  with  one  another 
and  with  the  results  of  the  platinum  chloride  method. 

By  using  10  grm.  of  soil  for  each  estimation  it  should  be  possible 
to  attain  a  higher  degree  of  accuracy. 

Potassium  in  Drushel  has  also  studied  the  application  of  the 
uri^f^oo^r  cobalti-nitrite  method  to  the  determination  of  potas- 
Lymph,  Milk,  sium  in  animal  fluids.*  Of  the  constituents  of  urine, 
ammonia  and  the  organic  substances,  especially  urea,  are  the 
only  ones  which  should  interfere  with  the  volumetric  method 
as  previously  described.  To  remove  these  without  the  loss  of 
potassium  is  apparently  the  only  new  problem  in  connection  with 
the  estimation  of  potassium  in  urine,  and  this  is  accomplished 
by  the  following  procedure:  Aliquot  portions  of  urine  of  10  to 
50  cm.3  each  are  measured  with  pipettes  or  a  burette  into  small 
platinum  evaporating  dishes,  and  evaporated  to  dryness  over 
the  steam  bath  in  a  good  draft  hood.  The  residues  are  best 
treated  by  acting  upon  them  with  5  cm.3  to  10  cm.3  of  a  9:  i 
nitric-sulphuric  acid  mixture  in  an  evaporating  dish  kept  covered 
until  the  first  violent  oxidation  is  over,  evaporation  to  dryness, 
and  ignition. 

By  this  treatment  the  ignition  of  the  residue  from  50  cm.3  of 
urine  may  be  readily  made  in  30  minutes  without  loss  of  mate- 
rial. The  residue  thus  prepared  is  treated  with  a  little  water 
and  a  few  drops  of  acetic  acid  to  dissolve  the  alkalies,  and  from 
this  point  the  process  is  carried  out  as  in  the  application  of  the 
cobalti-nitrite  method  to  the  estimation  of  potassium  in  pure 
salts,  as  previously  described. 

The  results  obtained  in  the  application  of  this  method  to  a 
number  of  specimens  of  human  urine  are  given  in  the  following 
table. 

*  W.  A.  Drushel,  Am.  Jour.  Sci.,  [4],  xxvi,  555. 


THE  ALKALI  METALS 


99 


Potassium  in  Urine. 


Urine 
taken. 

cm.3 

Specific 
gravity. 

Volume  in 
24  hours. 

cm.* 

K  found. 

K  in  34  hours, 
grm. 

Platinum 
chloride 
method.* 

grin. 

Cobalti- 
nitrite 
method. 

grm. 

10 
10 
IO 
IO 

25 
25 
25 

20 
20 

20 

20 
2O 
2O 
2O 

25 
50 
SO 
25 

1.025 

95° 

950 
QIO 
1130 

1500 

0.0293 
0.0292 

2.78 

2-77 
2.78 

2-77 
2.81 

2.84 

2.81 

3-44 
3-42 
3-47 

3-74 
3-74 
3-74 
3-74 

2-55 
2.52 
2-53 
2-54 

1.025 

0.0293 
0.0292 

0.0740 
o  .  0740 

0.0757 

0.0764 

o  .  0663 
0.0662 

0.0747 

1.025 
1.024 

0.0752 

o  .  0663 
0.0662 

I.OI8 

0.0425 

0.0839 

0.0843 

0.0424 

Modified  Lindo-Gladding  method,  after  removal  of  P,OS 


An  additional  difficulty  presents  itself  in  the  presence  of 
a  large  amount  of  protein  material  which  cannot  be  removed 
by  coagulation  and  nitration  without  a  considerable  loss  of 
potassium.  This  is  particularly  true  of  the  blood,  where  m<5st 
of  the  potassium  is  intimately  associated  with  the  protein 
of  the  corpuscles.  It  is  necessary  therefore  to  decompose 
protein  material  by  oxidation.  For  this  purpose  the  nitric- 
sulphuric  acid  mixture  works  less  satisfactorily  than  treatment 
with  concentrated  nitric  acid,  digestion  on*  the  steam  bath, 
dilution,  treatment  of  aliquot  portions  by  evaporation,  gentle 
ignition,  addition  of  sulphuric  acid  and  a  final  ignition;  or 
liquid  bromine  may  be  substituted  for  nitric  acid  in  the  first 
oxidation. 

Results  obtained  for  potassium  in  circulating  fluids  are  given 
in  the  following  table. 


100 


METHODS  IN  CHEMICAL  ANALYSIS 


Potassium  in  Blood  and  Lymph. 


K2O  found. 

Nature  of  fluid. 

Amount 
taken. 

grm. 

Platinum 
chloride 
method.* 

grm. 

Cobalti- 
nitrite 
method. 

grm. 

Per  cent. 

(-10.89 

0.0227 

0.0227 

0.21 

1   II  .  21 

o  .  0228 

o  20 

Defibrinated  pig's  bloodf. 

j  20.33 

0.0391 

0.19 

j  10.  16 

0.0203 

0.  21 

1  10-85 

O.O2II 

o  .  20 

L  11.03 



0.0236 

O.  21 

{30.00 

0.0174 

O.OI74 

0.058 

30  oo 

O    OI  7O 

o  060 

Sheep's  blood}  

o    •  ^^ 

Q 

w«  w*  /  y 

o  .  0181 

_         /: 

30.00 

O  .  OIoI 

O  .  OOO 

30.00 

0.0181 

0.060 

30.00 

0.0180 

0.060 

Serum  of  dog's  blood}  

(  IO.II 

<  10.04 

O.OO24 

0.0024 
0.0024 

0.024 
0.024 

(  10.07 

0.0023 

o  .  034 

f  10.28 

O.OOlS 

0.0018 

0.018 

10.01 



0.0019 

0.019 

Dog's  lymph} 

J   IO   OO 

O.OO2O 

O.O20 

j  10.03 

O.OOIQ 

O.OI9 

IO.  12 

O.OOI9 

O.OI9 

[10.32 

0.0022 

O.OO22 

O.O2I 

*  Modified  Lindo- Gladding  method,  after  removal  of  Ca,  Fe,  and  P2O6. 
t  Oxidized  by  bromine. 
J  Oxidized  by  nitric  acid. 


In  the  estimation  of  potassium  in  milk  suitable  amounts 
are  evaporated  to  dryness,  oxidized  with  concentrated  nitric 
acid,  again  evaporated  to  dryness,  and  ignited  gently  until 
nearly  all  organic  matter  has  been  burnt.  The  residue  is  mois- 
tened with  concentrated  sulphuric  acid  and  again  ignited.  In 
the  residue  thus  obtained  the  potassium  may  be  estimated 
by  the  cobalti-nitrite  method  as  applied  to  pure  salts  of  po- 
tassium. 

Results  obtained  by  the  method  in  the  analysis  of  cow's  milk 
are  given  on  the  following  page. 

The  modifications  introduced  by  Drushel  into  the  cobalti- 
nitrite  process  —  evaporation  nearly  to  dryness  and  oxidation 
of  the  nitrite  without  previous  removal  of  the  cobalt  —  add 


THE  ALKALI  METALS 


Tot 


greatly  to  its  usefulness.  The  necessity  of  long  standing  is 
avoided,  the  precipitate  may  be  filtered  and  washed  without 
trouble  and  the  manipulation  previous  to  titration  is  much 
shortened. 

Potassium  in  Milk. 


K2O  found. 

Milk  taken. 

Platinum 

Cobalti- 

chloride 
method.* 

nitrite 
method. 

Per  cent. 

grm. 

grm. 

grm. 

25-8 

0.0413 

o.  16 

25.8 

0.0432 

0.17 

25.8 

0.0428 

0.17 

51-6 

0.0833 

0.16 

25-7 

o  0454 

0.18 

25-7 

0.0457 

0.18 

2<    7 

o  04x1 

0.18 

*  Lindo-Gladding  method,  after  removal  of  Ca  and  P2O6. 


For  small  amounts  of  potassium  fairly  accurate  results  are 
obtained  by  using  the  permanganate  factor  calculated  from 
Adie  and  Wood's  formula  for  potassium  sodium  cobalti-nitrite. 
Sutton  has  suggested  that  more  accurate  results  may  be  secured 
by  obtaining  a  factor  empirically  from  a  pure  potassium  salt. 
The  results  recorded  above  were  obtained,  however,  by  using  the 
theoretical  factor  calculated  from  the  formula  of  Adie  and  Wood, 
K2NaCo(NO2)6.HoO,  their  analyses  of  potassium  sodium  cobalti- 
nitrite  having  been  verified*  by  the  analysis  of  a  carefully  pre- 
pared salt. 

The  chief  sources  of  error  in  the  method  appear  to  be  the  slight 
solubility  of  the  potassium  sodium  cobalti-nitrite,  one  part  in 
25,000  to  30,000  parts  of  water  at  room  temperature,  and  the 
tendency  of  the  precipitate  to  include  traces  of  sodium  cobalti- 
nitrite. 

The  method  requires  less  time  and  labor  than  the  chloroplati- 
nate  method,  and  is  applicable  in  the  presence  of  substances 
which  form  no  insoluble  cobalti-nitrites  and  which  neither  oxidize 
oxalic  acid  nor  reduce  potassium  permanganate. 


*  Am.  Jour.  Sci.,  [4],  xxvi,  562. 


'102  ;MEfHODS'IN  CHEMICAL  ANALYSIS 

RUBIDIUM  AND  CAESIUM. 
The  Spectroscopic  Determination  of  Rubidium. 

The  method  of  manipulation  previously  described  for  the 
Spectroscopic  determination  of  small  amounts  of  potassium  has 
been  adapted  by  Gooch  and  Phinney*  to  the  similar  determina- 
tion of  rubidium. 

In  the  work  upon  potassium  the  observations  of  the  red  line 
were  made  in  the  ordinary  laboratory  in  diffused  light,  but  pre- 
liminary experimentation  upon  the  rubidium  spectrum  immedi- 
ately developed  the  fact  that  the  blue  lines  are  better  to  work  by 
in  the  case  of  this  element,  and  that  a  dark  room  becomes  a 
necessity.  For  the  experiments  described  pure  rubidium  chloride 
was  prepared  by  many  fractional  precipitations  by  alcohol  out 
of  aqueous  solutions,  and  in  settling  the  question  as  to  the  coils 
which  should  be  used  the  choice  fell  upon  the  size  holding  0.02 
grm.  of  water  and  made  of  the  No.  28  wire,  the  superior  stiffness 
of  these  and  consequent  constancy  in  capacity  giving  them  the 
advantage  over  smaller  coils  of  finer  wire,  though  the  latter  are 
capable  of  bringing  out  greater  sensitiveness  of  the  reaction.  It 
was  found,  for  example,  that  under  the  most  favorable  conditions 
as  to  height  of  flame  and  width  of  slit,  0.0002  mg.  of  rubidium 
chloride  produced  the  blue  lines  at  the  last  limit  of  visibility 
when  the  larger  and  heavier  coil  was  in  the  flame;  with  a  coil 
holding  0.006  grm.  of  water  and  made  of  very  fine  wire  the  more 
immediate  volatilization  of  the  chloride  so  increased  the  delicacy 
of  the  Spectroscopic  reaction  that  it  was  possible  to  see  the  lines 
from  0.00005  mg.  of  the  salt.  These  figures  serve  as  an  indica- 
tion of  the  possible  delicacy  of  this  method  of  producing  spectra, 
but  it  should  be  remembered  that  all  eyes  do  not  see  the  rubidium 
lines  with  equal  ease. 

In  comparative  tests  of  brightness  it  was  found  best  to  employ 
as  the  standard  the  lines  given  by  amounts  of  the  chloride  not 
exceeding  0.0005  mg-  to  0.0007  mg->  to  set  the  slit  at  a  width  of 
0.2  mm.  and  to  bring  the  coils  to  the  flame  in  sets  of  three  — 
the  first,  usually  a  standard,  serving  to  fix  the  position  of  the 
lines  so  that  the  comparative  distinctness  of  the  lines  given  by 
the  other  two  might  be  the  more  readily  determined.  When 
that  dilution  had  been  found  at  which  the  test  was  barely  brighter 
*  F.  A.  Gooch  and  J.  I.  Phinney,  Am.  Jour.  Sci.,  [3],  xliv,  392. 


THE   ALKALI   METALS 


than  the  standard  and  that  dilution  at  which  the  test  was  barely 
weaker  than  the  standard,  it  was  assumed  that  the  mean  of  the 
numbers  of  cubic  centimeters  representing  these  two  volumes 
might  be  taken  as  the  volume  at  which  the  test  and  standard 
lines  were  equal.  The  amount  of  rubidium  in  the  test  solution 
was  then  calculated  by  multiplying  the  volume  in  cubic  centi- 
meters by  the  number  of  coilfuls  in  I  cm.3  and  the  product 
by  the  amount  of  rubidium  contained  in  a  coilful  of  the  stand- 
ard solution. 

The  results  of  two  experiments  with  pure  rubidium  chloride 
are  given  below. 

Determination  of  Rubidium  in  Pure  Rubidium  Chloride. 
Experiment  I. 


Standard. 
Rubidium  in  a 
coilful  (A  cm.*). 

Test  (known  to 
contain  10  mg.  Rb). 
Volume  in  cm.8 

Line  of  test 
compared  with 
standard. 

mg. 

0.0005 

340 

Brighter. 

0.0005 

370 

Equally  bright. 

0.0005 

370 

Brighter. 

0.0005 

39° 

Weaker. 

0.0005 

390 

Weaker. 

Found,  37°  *  39°  X  50  X  0.0005  =  9-5  nig. 

Taken 10.0  mg. 

Error 0.5  mg.  =  5  per  cent. 


Experiment  II. 


Standard. 

Rubidium  in  a 
coilful  (^j  cm.3). 

Test  (known  to 
contain  10  mg.  Rb). 
Volume  in  cm.3 

Line  of  test 
compared  with 
standard. 

mg. 

0.0005 

300 

Brighter. 

0.0005 

360 

Equally  bright. 

0.0005 

380 

Brighter. 

0.0005 

380          ' 

Brighter. 

0.0005 

390 

Brighter. 

0.0005 

400 

Weaker. 

0.0005 

410 

Weaker. 

Found,  39°  +  4°°  X  50  X  0.0005  =  9. 875 rag. 

Taken , 10  o 

Error o.  125  mg.  =  i .  25  per  cent. 


104  METHODS  IN  CHEMICAL  ANALYSIS 

These  results  make  it  plain  that  when  the  comparison  is  made 
between  solutions  of  pure  rubidium  chloride  the  spectroscopic 
method  is  capable  of  yielding  fair  approximations  to  truth. 
In  the  practical  determination  of  rubidium,  however,  the  ques- 
tion of  the  effect  of  the  presence  of  sodium  and  potassium  which 
naturally  accompany  it  is  of  importance. 

It  appears  from  practical  tests  that  within  limits  the  presence 
of  sodium  in  the  flame  increases  the  brilliance  of  the  rubidium 
spectrum.  The  brightness  of  the  lines  is  raised  under  the  con- 
ditions by  a  maximum  of  50  per  cent  by  the  presence  of  sodium 
up  to  40  per  cent  of  the  weight  of  the  rubidium,  and  increase 
in  the  amount  of  sodium  does  not  further  influence  the  bright- 
ness of  the  lines  until  the  proportion  of  sodium  to  rubidium  is 
as  ten  to  one;  or,  speaking  broadly,  the  dissociating  effect  of 
sodium  upon  the  rubidium  chloride  (to  which  the  brightening 
noted  is  to  be  attributed)  does  not  appear  to  be  materially 
different  whether  one  or  a  score  of  molecules  of  sodium  chlo- 
ride are  present  to  one  of  the  rubidium  chloride.  But  when  the 
proportion  of  sodium  to  rubidium  much  exceeds  ten  to  one 
the  glare  of  light  diffused  through  the  entire  spectrum  (though 
the  sodium  line  itself  may  be  cut  off)  begins  to  affect  the  vision, 
and  as  the  increase  advances  ultimately  extinguishes  the  rubidium 
lines. 

It  appears  also  that  potassium  chloride  produces  an  effect  simi- 
lar to  that  of  sodium  chloride,  the  brightness  of  the  rubidium  line 
increasing  by  a  maximum  of  50  per  cent  when  the  potassium  is 
present  to  between  two-thirds  and  twice  the  amount  of  the  rubid- 
ium; while  the  presence  of  potassium  in  the  proportion  five  to 
one  influences  the  visibility  unfavorably,  and  in  the  proportion 
of  thirty  to  one  extinguishes  the  rubidium  line  in  the  glare  of 
light.  It  is  necessary  therefore  either  to  effect  the  separation  of 
the  rubidium  from  sodium  and  potassium,  or  else  to  bring  test 
and  standard  to  the  same  condition  as  regards  the  presence  of 
these  elements,  before  any  reasonable  degree  of  accuracy  can  be 
expected  in  the  spectroscopic  determination  of  rubidium  as  it 
ordinarily  occurs  in  nature.  The  separation  from  sodium  is  easily 
accomplished  by  the  conversion  of  the  salts  to  the  form  of  chloro- 
platinates;  but  for  the  quantitative  separation  of  rubidium  from 
potassium  there  is  no  good  method  known.  The  practical  value 
of  the  spectroscopic  reaction  of  rubidium  for  purposes  of  quanti- 


THE  ALKALI  METALS  105 

tative  analysis  depends,  therefore,  upon  matching  potassium  lines 
as  well  as  the  rubidium  lines  (following  the  method  outlined  in 
the  determination  of  potassium  in  presence  of  sodium),  and  so 
bringing  the  lines  of  test  and  standard  equally  under  the  influence 
of  potassium.  It  has  been  shown  that  there  is  no  difficulty  in 
matching  solutions  of  potassium  by  means  of  the  red  line,  but 
the  convenience  of  using  the  spectroscope  without  readjust- 
ment throughout  an  entire  experiment  makes  a  comparison 
by  means  of  the  blue  line  highly  desirable  and  this  has  been 
found  to  be  feasible.  The  details  of  a  determination  of  rubid- 
ium in  presence  of  a  permissible  amount  of  potassium  are  given 
in  the  following  statement. 


Determination  of  Rubidium  in  Presence  of  Potassium. 


Standard   solution  containing 
(&cm.*), 

Test  solution  containing  8  mg.  rubidium  and  no  potassium. 


to  the  coilful 


Step  i. 

Step  2. 

Step  3. 

Step  4. 

Step  5. 

Preliminary  test 
for  Rb. 

Preliminary 
matching  of  K 
line. 

Rematching  of 
Rb  line. 

Readjustment 
of  K  line. 

Final  matching 
of  Rb  line. 

Test  at  20  cm.1 
gave  Rb  line 
like  standard. 

Test  at  20  cm.8 
gave  K  line  like 
standard  when 
i  mg.  of  K  had 
been  added. 

Test  at  35  cm.* 
gave  Rb  lines 
like  standard. 

Test  at  35  cm.* 
gave  K  line  like 
standard  when 
2  mg.  were  pres- 
ent. 

Test    at    35   cm.» 
gave  Rb  line  like 
standard. 

Found,  35X50X0.0005  =0.875  mg. 

Taken =0.8  mg. 

Error =0.075  mg.  =  g.4  per  cent. 

When  the  amount  of  potassium  present  is  so  great  as  to 
vitiate  the  test  for  rubidium  a  precipitation  by  alcohol  may  be 
utilized  to  remove  the  excessive  amount  of  the  potassium  salt. 
The  mixed  chlorides  are  dissolved  in  the  least  possible  quantity 
of  water  and  treated  with  absolute  alcohol;  the  precipitate  is 
filtered  off  and  washed  with  alcohol;  the  filtrate  and  washings 
are  evaporated  and  the  residue  dissolved  in  a  known  volume  of 
water  is  ready  for  the  spectroscopic  test.  Results  of  experi- 
ments conducted  in  this  manner  follow. 


io6 


METHODS  IN  CHEMICAL  ANALYSIS 


Rubidium  taken 
in  the  form  of 
chloride. 

Potassium  taken 
in  the  form  of 
chloride. 

Rubidium  found. 

Absolute  error. 

Percentage  error. 

mg. 

grm. 

mg. 

mg. 

per  cent. 

I 

O.I 

0.8 

O.2 

20 

2 

O.I 

i-7 

o-3 

15 

I 

O.  I 

0.9 

O.I 

IO 

The  error  of  the  process  is  manifestly  large,  and  only  roughly 
approximate  results  can  be  hoped  for  when  large  amounts  of 
rubidium  are  dealt  with;  but,  in  view  of  the  fact  that  the  only 
alternative  is  an  indirect  process,  even  this  great  error  may 
not  be  prohibitive  in  the  estimation  of  very  small  amounts  of 
rubidium.* 

The  Estimation  of  Caesium  and  Rubidium  as  the  Acid  Sulphates. 

Browning  f  has  shown  that  by  holding  the  temperature  between 
250°  and  270°  during  treatment  with  sulphuric  acid  suitable  salts 
of  caesium  and  rubidium  may  be  brought  with  a  fair  degree  of 
certainty  to  the  condition  of  the  acid  sulphates,  which  by  treat- 
ment with  ammonia  and  ignition  at  red  heat  yield  the  neutral 
sulphates.  Salts  of  potassium  and  sodium,  however,  when  heated 
at  the  same  range  of  temperature,  yield  pyrosulphates  reasonably 
stable  under  the  conditions.!  Lithium  salts  when  treated  simi- 
larly gave  no  evidence  of  the  existence  of  a  stable  acid  sulphate 
or  pyrosulphate.  Details  of  the  experiments  with  the  salts  of 
rubidium  and  caesium  are  given  below. 

A  weighed  amount  of  caesium  nitrate  was  placed  in  a  pre- 
viously weighed  platinum  crucible  and  treated  with  an  excess  of 
sulphuric  acid.  The  crucible  was  then  placed  upon  a  steam 
bath  until  the  water  and  nitric  acid  were  largely  expelled,  and 
then  removed  to  a  radiator,  consisting  of  a  porcelain  crucible 
fitted  with  a  pipe-stem  triangle  so  arranged  that  the  bottom  of 
the  platinum  crucible  was  about  midway  between  the  top  and 
bottom  of  the  porcelain  crucible.  This  improvised  radiator  was 
set  in  an  iron  ring  and  a  thermometer  placed  so  that  the  mer- 
cury bulb  was  on  a  level  with  the  bottom  and  close  to  the 
side  of  the  platinum  crucible.  An  ordinary  Bunsen  burner 

*  For  an  example  of  the  practical  use  of  this  method,  see  The  Excretion  of 
Rubidium,  Mendel  and  Slosson,  Am.  Jour.  Physiol.,  xvi,  152. 
t  Philip  E.  Browning,  Am.  Jour.  Sci.,  [4],  xii,  301. 
t  See  pages  79,  92. 


THE  ALKALI  METALS 


107 


served  as  the  source  of  heat  and  the  temperature  was  kept  so 
far  as  possible  between  250°  C.  and  270°  C.  After  the  fuming 
of  sulphuric  acid  had  ceased,  the  crucible  and  contents  were 
removed  to  a  desiccator,  cooled  and  weighed.  This  process  of 
heating  was  continued  for  half-hour  periods  until  the  weights 
were  constant.  The  results  given  show  that  by  regulating  the 
heat  at  a  temperature  between  250°  C.  and  270°  C.  caesium  may 

Caesium  Sulphates  by  Ignition. 


CsNO, 
taken. 

grm. 

CsHSO4 
calculated. 

grm. 

First 
constant 
weight. 

grm. 

Second 
constant 
weight. 

grm. 

Error  on 
CsHSO4. 

grm. 

Cs2S04 

calculated. 

grm. 

Cs2S04 
found. 

grm. 

Error  on 
Cs2S04. 

grm. 

O2OI  3 

O   2O?4 

O    2O2O 

-j-o  0007 

O   I7o6 

o  2013 

O   2OIO 

—  o  0003 

OIO32 

o  1217 

O   1  2OI 

—  o  0016 

0.1032 

0.1218 
o  1214 

O.I2I7 
0-1437 

o  143? 

0.1252 
0.1458 

o  1430 

0.1222 

+0.0005 

+0.0021 

—  o  0005 

0.0961 
0.1130 

o  .  0948 
o.  1118 

—0.0013 

—  O.OOI2 

o  1214 

O    143? 

o  1422 

—  o  0013 

OT  T  CQ 

o  1  3  ?6 

O    I  33O 

—  o  0026 

O    IO?6 

O    1  24? 

o  1272 

o  1248 

+o  0003 

o  1056 

O    1  24.? 

o  1  252 

+o  0007 

be  brought  with  a  fair  degree  of  certainty  to  the  condition  of  the 
acid  sulphate.  As  a  check  upon  the  results  the  acid  sulphate 
was,  in  a  few  cases,  treated  with  a  little  ammonium  hydroxide, 
the  excess  of  this  was  removed  upon  a  steam  bath  and  the  neutral 
sulphate  was  obtained  by  ignition  at  a  red  heat  to  a  constant 
weight.  These  determinations  agree  fairly  well  with  the  theory. 
The  same  procedure  was  followed  with  rubidium,  a  pure  ru- 
bidium chloride  having  been  chosen  as  the  starting  point.  The 
results  are  given  below.  No  tendency  was  observed  on  the  part 
of  these  elements  to  hold  sulphuric  acid  in  excess  of  the  amount 
necessary  for  the  formation  of  the  acid  sulphate. 

Rubidium  Sulphates  by  Ignition. 


RbCl 
taken. 

grm. 

RbHSO4 
calculated. 

grm. 

RbHSO4 

found. 

grm. 

Error, 
grm. 

Rb2S04 
calculated. 

grm. 

Rb2SO4 
found. 

grin. 

Error* 
grm. 

o.  1252 

o  1889 

o  1878 

—  O   OOII 

O.I2I2 
O.I23O 

o.  1829 
o  1856 

o.  1840 
o  1850 

+O.OOII 

—  o  0006 

0.1460 

0.1460 

6.OOOO 

0.1230 

0.1610 
0.1360 

0.1856 
0.2430 
0.2052 

0.1858 

0.2416 

0.2032 

+O.OOO2 
—  O.OOI4 
—  O.OO2O 

0.1357 
0.1777 
0.1501 

0.1350 
0.1772 
0.1490 

—0.0067 
—  0.0005 
—  O.OOII 

CHAPTER   III. 
COPPER;  SILVER;  GOLD. 

COPPER. 
The  Gravimetric  Determination  of  Copper  as  the  Sulphocyanate. 

As  early  as  1854  attention  was  drawn  by  Rivot*  to  the  pos- 
sibility of  estimating  copper  gravimetrically  by  weighing  as 
cuprous  sulphocyanate,  and  to  the  advantages  which  the  process 
afforded  in  separating  copper  from  other  metals.  Rivot's  pro- 
cedure consisted  in  dissolving  the  substance  to  be  analyzed  in 
hydrochloric  acid,  reducing  the  copper  with  hypophosphorous 
or  sulphurous  acid,  and  precipitating  with  potassium  sulpho- 
cyanate. The  precipitate,  dried  at  a  moderate  temperature, 
was  weighed  as  cuprous  sulphocyanate  and  then  as  a  control 
converted  by  ignition  with  sulphur  into  cuprous  sulphide  and 
weighed  in  that  condition.  In  spite  of  the  evident  advantages 
for  certain  purposes,  Rivot's  method,  in  its  original  form,  has 
never  come  into  general  use,  the  chief  reason  for  this  being 
apparently  the  difficulty  and  inaccuracy  attendant  upon  the 
weighing  of  the  precipitate  upon  dried  paper  filters,  a  process 
which  can  hardly  be  depended  upon  unless  managed  with  extreme 
care. 

Van  Namef  has  shown,  however,  that  the  process  is  accurate 
and  easily  managed  if  attention  is  given  to  the  necessary  con- 
ditions of  concentration  and  acidity,  and  the  precipitated  cuprous 
sulphocyanate  is  filtered  and  weighed  upon  asbestos  in  the  per- 
forated crucible. 

The  table  contains  results  obtained  as  follows:  A  suitable 
quantity  of  a  standard  copper  sulphate  solution  was  run  from  a 
burette,  diluted  to  a  convenient  volume,  a  few  cubic  centimeters 
of  a  concentrated  solution  of  ammonium  bisulphite!  added,  and 
the  copper  precipitated  by  an  excess  of  ammonium  sulpho- 

*  Compt.  rend.,  xxxviii,  868. 

f  R.  G.  Van  Name,  Am.  Jour.  Sci.,  [4],  x,  451;  xiii,  20. 
%  Prepared  by  saturating  strong  aqueous  ammonia  with  sulphur  dioxide. 

108 


COPPER;   SILVER;   GOLD 


109 


cyanate.  After  allowing  the  mixture  to  stand  for  a  few  minutes 
or  hours,  according  to  the  amount  of  free  acid  present,  the  pre- 
cipitate was  collected  upon  asbestos  in  a  weighed  crucible,  washed 
with  cold  water  and  dried  at  1 10°  until  no  further  loss  of  weight 
took  place. 

Copper  weighed  as  Cuprous  Sulphocyanate. 


Cu  taken. 

H2S04 
cone. 

HNH4SO3 
sat.  sol. 

NH4SCN 
approx. 
n/io. 

Final 
volume. 

Time  of 
standing. 

Cu  found. 

Error. 

grin* 

cm.1 

cm.8 

cm.8 

cm.8 

hours. 

gnu  . 

grm. 

0.0795 

None. 

5 

13 

68 

1 

0.0795 

0.0000 

0.0795 

None. 

3 

13 

66 

48 

0.0793 

—  O.OOO2 

0.0795 

None. 

3 

25 

78 

£ 

0.0796 

+O.OOOI 

0.0795 

None. 

3 

25 

78 

12 

0.0796 

+O.OOOI 

0.0795 

i-5 

10 

13 

85 

12 

0.0792 

—0.0003 

0.0795 

i-5 

8 

13 

105 

48 

0.0785 

—o.ooio 

0.0795 

i-5 

3 

25 

85 

4 

0.0783 

—  0.0012 

0.0795 

i-5 

5 

25 

85 

21 

0.0795 

0.0000 

0.0795 

S 

5 

25 

85 

3 

0.0797 

+O.OOO2 

HCl  cone. 

cm.1 

0.0795 

IO 

5 

25 

IOO 

20 

0.0795 

0.0000 

0.0795 

25 

IO 

25 

IOO 

28 

0.0784 

—  O.OOII 

Larger  amounts  of  copper  may  also  be  estimated  in  the  same 
way,  as  the  following  table  shows,  but  with  a  crucible  of  the 
ordinary  size  the  process  is  more  rapid  and  convenient  when  a 
smaller  weight  of  copper  is  taken. 

Copper  weighed  as  Cuprous  Sulphocyanate. 


Cu  taken. 

HjSO4  cone. 

NH4SCN 
approx.  w/io. 

Final  volume. 

Cu  found. 

Error. 

gnu. 

cm.s 

cm.8 

cm.8 

gnu  . 

grm. 

0.3175 
0.3I7S 
0.3175 

None. 
None. 
None. 

60 
60 
60 

500 
500 
500 

0.3176 

0.3177 
0.3176 

-f-o.oooi 

+O.OOO2 
+  O.OOOI 

0.3175 

10 

IOO 

500 

0.3175  . 

o.oooo 

HCl  cone. 

cm.8 

0.3175 

20 

IOO 

500 

0.3165 

—o.ooio 

In  solutions  containing  free  acid  the  precipitation  of  the  cop- 
per is  greatly  retarded,  and  the  mixture  should  be  allowed  to 
stand  for  several  hours,  or,  if  the  amount  of  acid  is  considerable, 
for  at  least  twenty-four  hours  before  filtering.  Precipitation  from 


110  METHODS  IN  CHEMICAL  ANALYSIS 

a  warm  solution  is  permissible,  but  boiling  the  liquid  causes  the 
precipitate  to  turn  brown  with  gradual  loss  in  weight,  and  is 
therefore  to  be  avoided. 

The  only  difficulty  which  is  likely  to  be  encountered  in  the  use 
of  this  method  is  a  tendency,  which  sometimes  appears,  for  traces 
of  the  precipitate  to  pass  through  the  filter  during  the  last  stages 
of  the  washing.  This  tendency  is  most  marked  with  precipitates 
from  concentrated  solutions  containing  little  or  no  free  acid.  It 
may  be  reduced  to  an  insignificant  amount  or  entirely  eliminated 
by  employing  one  or  more  of  the  following  expedients:  (i)  pre- 
cipitating in  dilute  solution ;  (2)  precipitating  in  the  presence  of 
free  acid ;  (3)  filtering  and  washing  under  light  pressure,  using 
a  rather  dense  (not  thick)  asbestos  mat;  (4)  washing  with  a 
decinormal  solution  of  ammonium  sulphocyanate.  The  last 
expedient  is  of  little  or  no  use  when  the  precipitate  is  to  be 
directly  weighed,  but  is  very  satisfactory  in  separating  copper 
from  other  substances.  Precipitation  in  acid  solution  is  the 
most  effective  method  of  obtaining  precipitates  which  are  easily 
filtered,  but  must  be  used  with  caution,  for  the  errors  from  incom- 
plete precipitation  may  easily  exceed  the  mechanical  losses  which 
the  acidity  was  employed  to  prevent. 

The  effect  of  hydrochloric  acid  of  various  concentrations  upon 
the  completeness  of  the  precipitation  was  studied  in  a  series  of 
experiments,*  in  which  the  principal  stress  was  laid  upon  deter- 
mining the  amounts  of  copper  left  in  solution  rather  than  the 
weights  of  the  precipitates.  The  procedure  was  as  follows: 
After  filtering  off  the  precipitate  the  copper  in  the  filtrate  was 
determined  by  evaporating  the  solution  with  nitric  acid  to  a 
small  bulk,  heating  in  a  platinum  crucible  over  a  radiator  to 
expel  sulphuric  acid  and  decompose  interfering  substances,  dis- 
solving the  residue  in  nitric  acid,  filtering,  electrolyzing  and 
weighing  the  copper.  The  electrolytic  deposit  was  then  redis- 
solved  and  the  copper  estimated  more  accurately  by  a  colori- 
metric  method  based  on  comparison  of  the  intensity  of  the 
brown  color  produced  upon  the  sample  by  potassium  ferrocyanide 
with  that  of  a  variable  standard  of  known  copper  content. 

The  results  of  this  series  of  experiments  may  be  summarized 
as  follows :  Allowance  must  be  made  for  the  amount  of  hydro- 
chloric acid  used  up  by  interaction  with  the  ammonium  bisulphite 
*  Am.  Jour.  Sci.,  [4],  xiii,  20. 


COPPER;   SILVER;    GOLD 


III 


solution  forming  ammonium  chloride  and  sulphur  dioxide.  One 
cubic  centimeter  of  a  bisulphate  solution,  prepared  by  saturating 
strong  aqueous  ammonia  with  sulphur  dioxide,  may  neutralize 
by  this  reaction  about  nine-tenths  of  a  cubic  centimeter  of  hydro- 
chloric acid,  sp.  gr.  1.18.  In  order  that  the  amount  of  copper 
left  in  solution  may  not  exceed  o.i  mg.  per  100  cm.3  of  nitrate, 
the  concentration  of  effective  hydrochloric  acid,  i.e.,  that  remain- 
ing after  interaction  with  the  bisulphite,  stated  in  volume  per 
cent  of  the  concentrated  acid  (cubic  centimeters  of  acid  of  sp.  gr. 
1.18  for  100  cm.3  final  volume  of  solution),  should  not  be  above 
0.8  when'  the  excess  of  sulphocyanate  employed  is  small  (20  per 
cent  above  the  theory),  but  may  be  as  high  as  3  per  cent  if  ten 
times  the  theoretical  amount  of  sulphocyanate  be  employed.  A 
suitable  degree  of  acidity  for  precipitating  copper  under  ordinary 
conditions  is  given  by  0.5  to  i.o  per  cent  of  hydrochloric  acid, 
expressed  as  above,  using  from  five  to  ten  times  the  theoretical 
amount  of  sulphocyanate. 

As  far  as  could  be  judged  from  a  limited  number  of  determina- 
tions made  in  the  presence  of  sulphuric  acid,  the  above  holds  true 
for  the  equivalent  amount  of  sulphuric  acid. 


Cu,(SCN), 
taken. 

Volume  of 
liquid. 

cm.3 

HC1  (sp.  gr.  1.18). 
cm.» 

NH4SCN. 

grni. 

Cu  in  filtrate, 
grm. 

0-3 

2OO 

0.06035 

0-3 

2OO 

0.00040 

0.25 

2OO 
2OO 

IO. 

1.52* 

0.0050 
0.00018 

o-3 

2OO 

. 

0.767 

0.00007 

o-3 

2OO 

o.  19! 

0.00004 

0-3 

2OO 

2 

.... 

0.0019 

0-3 
0-3 

2OO 
200 
200 

2 
2 
2 

2.5 

1-77** 
0.19 

0.0013 
0.0009 
0.0006 

NH4C1. 

grm. 

0-3 

200 

IO 

0.0031 

0-3 
0-3 

200 

200 

IO 
•ft 

0.19 

0.00013 
0.00045 

0-3 

2OO 

I 

0.19 

0.00005 

*  Solution  w/io  In  respect  to  NH4SCN. 

t  Solution  w/2o  in  respect  to  NH4SCN. 

J  Solution  w/8o  in  respect  to  NH4SCN. 
**  HC1  and  NH4SCN  present  in  equivalent  amounts. 
ft  Solution  approximately  w/io  in  respect  to  NH4C1. 


112  METHODS  IN  CHEMICAL  ANALYSIS 

The  effect  of  varying  concentrations  of  different  reagents  in- 
volved in  the  process  upon  the  solubility  of  cuprous  sulpho- 
cyanate  is  shown  in  a  rough  way  by  the  preceding  table.  As 
no  stirring  was  employed  the  figures  have  no  absolute  value,  but 
serve  merely  to  give  an  idea  of  the  relative  magnitude  of  the 
solubilities  in  question. 

Weighed  amounts  of  cuprous  sulphocyanate  prepared  by  pre- 
cipitation in  the  usual  way,  thoroughly  washed,  and  dried  at 
105°,  were  allowed  to  stand  in  the  solutions  to  be  tested  from 
40  to  50  hours.  After  careful  filtering  through  asbestos  the 
copper  in  the  clear  filtrate  was  estimated  by  electrolysis,  or,  in 
cases  where  the  amount  was  small,  by  the  colorimetric  method 
referred  to  above. 

The  solubility  in  presence  of  either  hydrochloric  acid,  ammo- 
nium chloride  or  a  large  amount  of  ammonium  sulphocyanate 
is  considerable.  It  is  lowest  in  dilute  solutions  of  ammonium 
sulphocyanate,  and  the  presence  of  a  small  amount  of  this  salt 
lessens  the  solubility  in  hydrochloric  acid,  and  in  solutions  of 
ammonium  chloride. 

Separation  of  From  the  nature  of  the  process  it  is  evident  that 
Copper  from  jt  js  mucn  iess  likely  to  be  interfered  with  by  the 

Bismuth,  Anti-  J  .  . 

mony,  Tin  and  presence  of  other  metals  than  the  other  gravimetric 
Arsenic.  methods  for  copper,  and  may,  therefore,  be  directly 

applied  with  good  results  in  many  cases  where  the  use  of  the 
electrolytic  or  the  oxide  method  would  involve  a  previous  sepa- 
ration. Van  Name*  has  tested  the  method  for  the  separation 
of  copper  from  bismuth,  antimony,  tin  and  arsenic. 

Having  copper  present  with  these  metals  in  a  solution  con- 
taining free  hydrochloric  acid,  tartaric  acid  was  added  to  aid 
in  preventing  the  formation  of  insoluble  products  of  hydrolytic 
action,  and  the  copper  then  precipitated  as  cuprous  sulphocya- 
nate in  the  usual  way. 

All  the  determinations  were  allowed  to  stand  fifteen  hours 
or  more  before  filtering  to  insure  completeness  of  precipitation. 
The  filtering  was  performed  upon  asbestos  in  a  perforated 
crucible.  The  precipitate  was  thoroughly  washed  with  cold  water 
and  dried  at  105°  to  a  constant  weight. 

In  the  following  table  are  results  obtained  by  this  procedure. 
The  acidity  was  kept  within  the  limits  shown  above  to  be  safe, 
*  Am.  Jour.  Sci.,  [4],  xiii,  138. 


COPPER;   SILVER;   GOLD 


and  the  amount  of  sulphocyanate  used  was  in  most  cases  about 
ten  times  the  theory. 

Copper  in  Presence  of  Bismuth,  Tin,  Antimony,  and  Arsenic. 
Final  Volume  200  cm.3. 


Cu. 
taken. 

grm. 

Bi. 

grm. 

HC1 

(sp.  gr. 
about 
I.I7). 

cm.» 

Tartar!  c 
acid. 

grm. 

HNH4SO3 
sat.  sol. 

cm.1 

NH4SCN 
approx. 

W/IO. 

cm.s 

Cur 

(SCN)2 
found. 

grm. 

Calcu- 
lated 
asCu. 

grm. 

Error, 
grm. 

0.0793 

O.2 

6 

2 

60 

0.1504 

0.0786 

—  0.0007 

0.0793 

O.I 

6 

2 

IOO 

0.1512 

o  .  0790 

-0.0003 

0.0793 

0-3 

6 

2 

125 

O.I5I5 

0.0792 

—  o.oooi 

0.0793 

O.2 

6 

2 

125 

0.1518 

0.0793 

o.oooo 

0.0793 

O.2 

6 

2 

125 

0.1519 

0.0794 

+O.OOOI 

0.0793* 

0.2 

6 

2 

230 

0.1519 

0.0794 

+O.OOOI 

Sn  taken  as 

SnCU+HCl. 

grm. 

0.0793 

O.2 

5 

I 

2 

40 

0.1502 

0.0785 

—0.0008 

0.0793 

O.2 

6 

I 

2 

125 

0.1514 

0.0791 

—  O.OOO2 

0.0793 

O.2 

5 

I 

2 

130 

0.1516 

0.0792 

—o.oooi 

Taken  as 

SnCl,+HCl. 

grm. 

0.0793 

0.2 

6 

I 

2 

125 

0.1529 

0.0799 

+0.0006 

As. 

grm. 

0.0793 

O.  2 

6 

I 

2 

125 

0.1523 

0.0796 

+0.0003 

Sb. 

grm. 

0.0793 

O.2 

6 

2 

2 

125 

0.1518 

0.0793 

0.0000 

As,  Bi,  Sb, 

Sn  of  each. 

grm. 

0.0795 

O.  I 

6 

2 

2 

130 

0.1523 

0.0796 

+0.0001 

0.0795 

O.I 

6 

2 

2 

130 

0.1525 

0.0797 

+0.0002 

*  Final  volume  300  cm.  * 

If  bismuth  is  present  in  considerable  amount,  a  good  deal  of 
hydrochloric  acid  is  needed,  and  there  is  danger  that  interaction 
with  the  precipitants  may  reduce  the  acidity  to  the  point  where 
hydrolysis  and  precipitation  of  the  bismuth  begins.  In  such 
cases  preliminary  blank  tests  must  be  carried  out  to  determine 
the  minimum  concentration  of  hydrochloric  acid  which  may  be 
employed  under  the  conditions.  With  antimony  the  effective- 
ness of  the  tartaric  acid  is  so  great  that  this  difficulty  does  not 
arise  if  enough  tartaric  acid  is  used.  Tin  in  the  stannous  con- 
dition sometimes  forms  a  slight  precipitate  of  sulphur  on  stand- 


114  METHODS  IN  CHEMICAL  ANALYSIS 

ing  in  contact  with  the  bisulphite,  and  it  is,  therefore,  advisable 
to  oxidize  it  at  the  outset  to  the  stannic  state. 

It  is  evidently  possible  to  estimate  copper  by  this  method  in 
the  presence  of  bismuth,  antimony,  tin  and  arsenic,  either 
separately  or  in  any  combination.  To  separate  copper  from 
unknown  quantities  of  bismuth,  or  from  mixtures  containing 
bismuth,  the  following  procedure  is  recommended:  Having  the 
copper  and  bismuth  in  hydrochloric  acid  solution,  add  tartaric 
acid,  and,  after  diluting  if  necessary,  determine  by  blank  tests 
with  small  aliquot  portions  of  the  solution  how  much  ammonium 
bisulphite  can  be  added  to  the  whole  without  precipitating  the 
bismuth.  Then,  keeping  the  bisulphite  well  within  this  limit, 
carry  out  the  precipitation  of  the  copper  as  already  described, 
using  a  considerable  excess  of  ammonium  sulphocyanate.  Where 
bismuth  is  absent,  antimony  and  tin  may  be  treated  in  the  same 
way,  but  the  latitude  possible  in  the  adjustment  of  the  conditions 
is  so  much  greater  with  these  metals  that  preliminary  tests  will 
seldom  be  needed.  For  the  separation  from  arsenic  no  special 
precautions  are  required. 

The  Determination  of  Copper  as  Cuprous  Iodide  and  Separation 
from  Cadmium. 

The  separation  of  copper  from  cadmium  by  the  precipitation 
of  the  cuprous  iodide  by  appropriate  means  has  long  been  known. 
Pisani*  mentions  the  fact  that  potassium  iodide  can  be  used  to 
effect  precipitations,  and  claims  that  a  satisfactory  separation 
can  be  made  in  this  way.  Flajolotf,  stating  that  potassium 
iodide  cannot  be  used  as  a  precipitant  on  account  of  the  solu- 
bility of  cuprous  iodide  in  that  reagent,  and  that  hydriodic  acid 
cannot  be  employed  if  nitric  acid  is  present,  recommends  that 
the  solution  containing  copper  be  brought  to  acidity  with  sul- 
phuric acid,  that  a  considerable  excess  of  sulphurous  acid  be 
added,  and  that  the  precipitation  be  effected  by  hydriodic  acid. 
Kohnerf  reviews  the  various  methods  for  the  separation  of 
copper  from  cadmium  and  states  that  the  iodide  method  is 
impracticable  on  account  of  the  solubility  of  cuprous  iodide  both 
in  excess  of  hydriodic  acid  and  in  potassium  iodide. 

*  Compt.  rend.,  xlvii,  294. 

t  Jour,  prakt.  Chem.,  Ixi,  105. 

J  Zeit.  anal.  Chem.,  xxvii,  203;  Jour.  Anal.  Chem.,  iii,  339. 


COPPER;   SILVER;   GOLD 


Browning*  has  fixed  conditions  under  which  accurate  results 
may  be  obtained  by  this  method. 

According  to  the  best  procedure  shown,  the  sulphates  of  copper 
and  cadmium,  in  amount  not  exceeding  0.25  grm.  of  each  metal, 
are  dissolved  in  25  cm.3  of  water  and  treated  with  I  grm.  or  2 
grm.  of  potassium  iodide,  f  The  mixture  is  evaporated  to  dry- 
ness  to  expel  iodine  and  then  treated  with  100  cm.3  of  water. 
Filtration  is  made  under  gentle  pressure  upon  the  asbestos  felt 
in  the  perforated  crucible.  It  is  advisable,  on  account  of  the 
tendency  of  cuprous  iodide  to  pass  through  the  filter,  to  use  a 
fairly  thick  felt  and  to  keep  it  moist  and  under  pressure  during 
the  filtration.  The  precipitate  is  washed  thoroughly  with  either 
hot  or  cold  water,  dried  at  I2O°-I5O°  and  weighed  as  cuprous 
iodide. 

Copper  Weighed  as  Cuprous  Iodide. 


Copper  taken, 
grm. 

Copper  found 
(weighed  as  C\i2I2) 

grm. 

Error  on  copper, 
grm. 

KI  used, 
grm. 

Final  volume  of 
liquid. 

cm.1 

0.1194 

0.1196 

+O.OOO2 

I 

IOO 

0.1191 

o.  1194 

+0.0003 

I 

IOO 

0.1193 

O.II93 

O.OOOO 

2 

IOO 

0.0049 

0.0045 

—  O.OO04 

2 

IOO 

0.0051 

0.0047 

—  O.0004 

2 

IOO 

0.1195 

O.II95 

0.0000 

3 

IOO 

0.1192 

0.1188 

—0.0004 

4 

IOO 

Copper  Separated  from  Cadmium  and  Weighed  as  Cuprous  Iodide. 


Copper 

taken. 

Copper 

found 
(weighed 
as  Cu2I2). 

Error  on 
copper. 

Cadmium 
taken. 

Cadmium 
found 
(weighed 
as  CdO). 

Error  on 
cadmium. 

.:'! 
KI  used. 

Found 
volume 
of  liquid. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

cm.» 

0.2383 

0.2386 

+0.0003 

O  .  0484 

o  .  0490 

+0.0006 

2 

IOO 

o.  1192 

o.  1185 

—  0.0007 

0.2439 

0.2430 

+O.OOOI 

2 

IOO 

0.1193 

0.1194 

+O.OOOI 

0.1942 

0.1942 

O.OOOO 

2 

IOO 

O.  I2OI 

O.  I2OI 

o  .  oooo 

0.2426 

o.  2428 

+O.OOO2 

2 

IOO 

O.II93 

0.1193 

O.OOOO 

0.2436 

0.2433 

—0.0003 

2 

IOO 

0.0239 

0.0238 

—  O.OOOI 

0.1934 

0.1932 

—  O.OOO2 

I 

IOO 

0.0236 

0.0239 

+0.0003 

o.  1942 

0.1936 

—  0.0006 

I 

IOO 

0.0239 

0.0242 

+0.0003 

0.1444 

o.  1442 

—  0.0002 

I 

IOO 

0.0238 

0.0238 

0.0000 

0.1467 

0.1461 

—  O.OOO6 

I 

IOO 

1 

*  Philip  E.  Browning,  Am.  Jour.  Sc.,  [3],  xlvi,  280. 

t  For  the  results  of  study  of  the  effect  of  excess  of  potassium  iodide,  free 
acid  and  concentration  upon  the  solubility  of  cuprous  iodide,  see  page  121. 


Il6  METHODS  IN  CHEMICAL  ANALYSIS    . 

The  table  contains  results  of  experiments  made  according  to 
this  procedure  with  pure  copper  sulphate,  and  with  copper  sul- 
phate and  cadmium  sulphate  in  mixture. 


The  Electrolytic  Determination  of  Copper. 

Gooch  and  Medway*  have  applied  the  rotating  cathode  to  the 
rapid  electrolytic  determination  of  copper,  making  use  of  the 
apparatus  shown  in  Fig.  I3.f 

The  deposition  of  copper  from  a  solution  of  the  sulphate  was 
first  attempted,  and  the  procedure  was  as  follows:  The  solution, 
50  cm.3  in  volume,  was  placed  in  a  i5O-cm.3  beaker  and  acidulated 
to  give  better  conductivity.  The  stand  carrying  the  beaker  was 
raised  until  the  liquid  covered  about  two-thirds  of  the  crucible 
adjusted  to  the  shaft,  thus  giving  a  cathode  surface  of  about 
30  cm.2  The  anode  was  introduced  and  the  motor  started. 
The  wires  from  the  storage  batteries  were  connected  and  the  cur- 
rent was  allowed  to  pass  through  the  solution.  The  duration  of 
the  electrolysis  was  varied  according  to  the  strength  of  current 
used;  but  in  each  case,  after  the  deposit  was  nearly  complete, 
the  current  from  the  batteries  was  shut  off,  the  motor  stopped, 
and  the  sides  of  the  beaker,  the  platinum  anode  and  the  crucible 
were  carefully  washed  with  a  fine  jet  of  water,  the  motor  was 
again  started  and  the  current  allowed  to  pass  for  the  remaining 
time. 

When  the  deposit  was  complete  the  crucible  was  removed  and 
washed,  first  with  water,  then  with  alcohol,  and  finally  was  dried 
by  passing  it  over  a  flame. 

Sulphuric  acid  (6  or  7  drops  of  the  dilute  acid  —  1:3)  was 
generally  used  to  acidulate  the  solution,  since  it  was  found  that 
the  copper  was  deposited  in  less  time  with  sulphuric  acid  than 
with  nitric  acid  present.  Experiments  in  which  small  amounts 
of  nitric  acid  (6  to  9  drops  of  the  dilute  acid  — 1:3)  were  used 
show  that  the  copper  may  also  be  deposited  completely  in  pres- 
ence of  this  acid. 

The  following  tables  show  the  results  of  a  series  of  experiments 
made  as  described.  The  standard  of  the  solution  of  copper  sul- 
phate was  fixed  by  the  usual  slow  method  of  electrolytic  analysis. 

*  F.  A.  Gooch  and  H.  E.  Medway,  Am.  Jour.  Sci.,  [4],  xv,  320. 
t  See  page  12. 


COPPER;    SILVER;    GOLD 


Solution  of  CuSOt  Acidulated  with 


Copper  taken, 
grm. 

Copper  found, 
grm. 

Error, 
grm. 

Current. 

amp. 

N.  D.  100 

Time, 
min. 

0.0651 

0.0652 

+0.0001 

0.8 

2.7 

25 

0.0651 

0.0652 

+0.0001 

0.8 

2.7 

15 

0.0651 

0.0651 

0.0000 

i 

3-3 

10 

0.0651 

o  .  0649 

—  O.OOO2 

i 

3-3 

IO 

0.0651 

o  .  0648 

—0.0003 

i 

3-3 

IO 

o.  1272 

0.1272 

o.oooo 

2-5 

8-3 

15 

o.  1272 

0.1271 

—  O.OOOI 

2-5 

8-3 

15 

0.1272 

o.  1271 

—  O.OOOI 

2-5 

8-3 

IS 

0.1272 

0.1270 

—  O.OOO2 

3 

10 

13 

o.  1272 

0.1268 

—0.0004 

3 

10 

12 

0.2548 

0.2548 

0.0000 

3 

IO 

2O 

0.2548 

0.2548 

0.0000 

4 

13-3 

2O 

0.2548 

0.2550 

+O.OOO2 

4 

13.3 

2O 

0.2548 

0.2546 

—  O.OOO2 

4 

!3-3 

15 

0.2548 

0.2545 

—0.0003 

4 

13-3 

15 

Solution  of  CuSO^  Acidulated  with  HN03. 


Copper  taken. 
grm. 

Copper  found, 
grm. 

Error, 
grm. 

Current, 
amp. 

N.  D.  m 

Time, 
min. 

0.0651 

0  .  0648 

—  0.0003 

I 

3-3 

35 

0.0651 

0.0652 

+O.OOOI 

0.8 

2-7 

30 

0.0651 

o  .  0650 

—  O.OOOI 

0.8 

2-7 

2S 

0.0651 

o  .  0649 

—  O.OOO2 

i 

3-3 

25 

0.0651 

0.0650 

—  O.OOOI 

0.8 

2-7 

25 

0.0651 

0.0652 

+O.OOOI 

i 

3-3 

35 

0.0651 

o  .  0648 

—  0.0003 

i 

3-3 

30 

0.0651 

o  .  0650 

—  O.OOOI 

i-5 

5 

25 

0.0651 

0.0650 

—  O.OOOI 

!-5 

5 

25 

0.0651 

0.0647 

—0.0004 

1.8 

6 

20 

It  has  been  shown  also  by  Medway*  that  a  crucible  of  silver 
may  be  substituted  for  the  platinum  crucible  as  the  rotating 
cathode,  in  the  deposition  of  copper  from  the  acidulated  sulphate 
solution,  with  results  that  leave  little  to  be  desired  on  the  score 

of  accuracy. 

Deposition  upon  the  Silver  Crucible. 


Copper  taken, 
grm. 

Copper  found, 
gnu. 

Error, 
grm. 

Current, 
amp. 

N.  D.  m 

Time, 
min. 

0.1088 

0.1086 

—  0.0002 

2 

6.6 

15 

o.ioss 

o.  1090 

+0.0002 

2 

6.6 

15 

0.1088 

0.1084 

—  0.0004 

1-5 

5 

15 

0.1088 

0.1085 

—  0.0003 

2 

6.6 

is 

0.1088 

o.  1080 

—  O.OOOS 

2 

6.6 

15 

0.1041 

o.  1041 

O.OOOO 

2 

6.6 

15 

o.  1041 

o.  1046 

+0.0005 

2 

6.6 

15 

0.1041 

0.1039 

—  0.0002 

2 

6.6 

15 

Am.  Jour.  Sci.,  [4],  xviii,  180. 


Il8  METHODS  IN  CHEMICAL  ANALYSIS 

To  remove  the  copper  from  the  crucible,  the  deposit  is  rubbed 
off  as  much  as  possible  and  the  rest  may  be  dissolved  in  a  strong 
boiling  solution  of  hydrochloric  acid  with  but  trifling  loss  of 
silver,  as  is  shown  in  the  statement  of  the  results  of  two  experi- 
ments given  below: 


I. 

II. 

Weight  of  crucible  before  treatment 

•?  6  0080 

36  0062 

Weight  of  crucible  after  treatment  

36.0x562 

36  0041 

Loss  of  silver  

o  .  002  7 

O.OO2I 

It  appears  that  the  silver  crucible  may,  with  some  economy 
and  without  sacrifice  of  accuracy,  be  substituted  for  the  platinum 
crucible  used  as  a  rotating  cathode  in  the  electrolytic  determi- 
nation of  copper. 

The  lodometric  Estimation  of  Copper. 

When  potassium  iodide  is  added  to  a  suitable  solution  of  a 
cupric  salt,  cuprous  iodide  is  precipitated,  while  iodine  equiva- 
lent to  the  amount  of  iodine  fixed  in  the  cuprous  iodide  is  liber- 
ated. This  reaction  has  been  made  the  basis  of  an  iodometric 
method  for  the  determination  of  copper,  the  first  suggestion  of 
such  a  method  having  apparently  been  made  by  De  Haen  in 
1854.  In  this  process  cupric  sulphate  was  treated  in  solution 
with  potassium  iodide  and  the  free  iodine  determined  by  sul- 
phurous acid  according  to  Bunsen.  From  the  amount  of  iodine 
thus  found  the  copper  was  calculated,  according  to  the  equation 

2  CuSO4  +  4  KI  =  2  K2SO4  +  Cu2I2  +  I2. 

This  method  was  mentioned  in  the  following  year  by  Mohr,* 
with  the  modification  suggested  by  Schwarz  that  the  free  iodine 
be  determined  by  sodium  thiosulphate  instead  of  by  sulphurous 
acid.  E.  O.  Brown, f  apparently  without  knowledge  of  De  Haen's 
previous  work,  proposed,  in  1857,  similar  procedure,  and  in  1868 
the  method  with  slight  modification  was  presented  again  by 
Rtimpler.t  Concerning  the  utility  of  the  method  opinions  have 

*  Titrirmethode,  page  387. 

t  Jour.  Chem.  Soc.,  x,  65. 

}  Jour,  prakt.  Chem.,  cv,  193. 


COPPER;    SILVER;    GOLD 

varied.  Mohr  never  favored  it.  As  late  as  1877,  Mohr,*  after 
quoting  Meidinger  to  the  effect  that  cuprous  iodide  freshly  pre- 
cipitated and  washed  is  capable  of  taking  up  iodine,  and  Carl 
Mohr's  criticism  that  potassium  iodide  acts  upon  cuprous  iodide 
according  to  the  concentration,  states  that  the  method  is  not 
exact  and  has  nowhere  found  practical  application.  On  the 
other  hand,  Freseniusf  recommends  the  method  for  the  deter- 
mination of  small  amounts  of  copper,  noting  that  ferric  salts 
and  other  substances  capable  of  setting  free  iodine  from  an  acidi- 
fied solution  of  potassium  iodide  must  not  be  present,  and  indi- 
cates the  most  favorable  procedure.  The  copper  salt  treated, 
he  says,  should  be  the  sulphate,  preferably  in  neutral  solution, 
though  a  moderate  amount  of  sulphuric  acid  is  not  objectionable. 
Much  free  sulphuric  acid  and  all  free  nitric  acid  should  be  neu- 
tralized by  sodium  carbonate,  and  the  precipitate  dissolved  in 
acetic  acid,  an  excess  of  which  does  no  harm  in  the  iodometric 
process. 

Of  recent  writers,  some  have  favored  the  method,  while  others 
have  commented  upon  it  unfavorably.  Low  t  has  been  out- 
spoken in  praise,  to  the  extent  of  declaring  a  preference  for  this 
method  in  the  most  accurate  technical  work  over  all  other 
methods,  even  the  electrolytic  method. 

According  to  Low's  earlier  modification,  metallic  copper  is 
dissolved  in  nitric  acid,  the  solution  is  freed  from  nitrogen  oxides 
by  boiling,  a  considerable  amount  of  zinc  acetate  is  added,  and 
in  the  solution  having  a  volume  of  50  cm.3  an  excess  of  solid 
potassium  iodide  is  dissolved.  Zinc  acetate  is  preferred  to  sodium 
acetate  to  take  up  the  free  nitric  acid.  It  is  said  that  an  excess 
of  potassium  iodide  is  necessary  to  insure  rapidity  of  action  and 
is  harmless.  According  to  the  later  modification  of  this  method , 
Low  prepares  the  cupric  salt  by  dissolving  the  metal  in  nitric 
acid  (sp.  gr.  about  1.20),  boils  the  solution,  adds  bromine  water 
to  destroy  the  nitrogen  oxides,  boils  to  expel  the  bromine,  treats 
with  ammonium  hydroxide  in  excess,  adds  acetic  acid  and  boils 
again  if  necessary  to  get  a  clear  solution.  The  advantage  of 
using  an  excess  of  potassium  iodide  is  emphasized,  and  the  state- 
ment is  made  that  unless  an  excess  of  this  reagent  is  present  the 

*  Titrirmethode,  5  Aufl.,  288. 

t  Quant.  Anal.,  6te  Aufl.,  335,  1875. 

|  Jour.  Am.  Chem.  Soc.,  18,  468;  24,  1083, 


120  METHODS  IN  CHEMICAL  ANALYSIS      . 

reaction  does  not  proceed  to  completion  until  the  titration  of  the 
free  iodine  takes  place.  Low  recommends  the  use  of  I  grm.  of 
potassium  iodide,  an  excess  of  0.6  grm.,  for  every  0.075  grm.  of 
copper. 

As  a  result  of  elaborate  study,  Moser*  has  reached  the  con- 
clusion that  the  reaction  by  which  cuprous  iodide  is  formed 
from  potassium  iodide  and  cupric  sulphate  in  neutral  solution 
is  complete  at  very  high  concentration  of  the  solution;  that  the 
completeness  of  the  reaction  is  greatly  affected  by  the  volume  of 
liquid ;  that  the  amount  of  potassium  iodide  employed  is  almost 
without  influence  either  in  neutral  solution  or  in  acid  solution; 
and  that  the  presence  of  free  sulphuric  acid  even  in  large  amounts 
or  of  hydrochloric  acid  present  in  amount  equivalent  to  the 
cupric  sulphate  is  advantageous.  Moser  recommends,  there- 
fore, the  addition  of  sulphuric  acid  for  the  purpose  of  bringing 
the  reaction  to  completion. 

According  to  Fernekes  and  Koch,f  an  excess  of  acetic  acid  does 
not  influence  titrations,  while  a  certain  amount  of  potassium 
iodide —  1.5  grm.  to  2  grm.  for  0.0038  grm.  of  copper,  and  2.5 
grm.  for  0.0939  grm.  of  copper  —  must  be  added  to  bring  about 
complete  action  in  a  volume  of  100  cm.3 

Quite  recently  Cantoni  and  RosensteinJ  have  tested  the  reac- 
tion between  potassium  iodide  and  a  cupric  salt  under  various 
conditions ;  but  these  investigators  do  not  give  the  absolute  val- 
ues of  the  amounts  of  copper  taken  and  found,  merely  recording 
the  relative  effects  of  varying  conditions.  From  the  record  of 
their  results  it  would  appear  that  a  fivefold  increase  of  the  mini- 
mum amount  of  potassium  iodide  added  to  portions  of  100  cm.3 
of  solution  containing  the  same  amount  of  copper  salt  is  without 
influence  upon  the  result;  that  increase  of  volume  from  loocm.3 
to  350  cm.3,  other  conditions  being  the  same,  may  affect  the 
results  by  as  much  as  5  per  cent  of  their  value.  The  authors 
conclude  that  the  method  gives  good  results  under  properly  con- 
trolled conditions. 

So  evidence  and  opinions  as  to  the  effect  of  various  conditions 
in  the  process  are  contradictory,  the  chief  matters  of  difference 
being  the  influence  of  an  excess  of  potassium  iodide  used  as  the 

*  Zeit.  anal.  Chem.,  xliii,  597. 

f  Jour.  Am.  Chem.  Soc.,  xxvii,  1229. 

j  Bull.  Soc.  Chim.,  [3],  xxxv,  1067-73. 


COPPER;    SILVER;    GOLD  121 

precipitant,  the  dilution  at  which  the  precipitation  should  take 
place,  and  the  effect  of  acids  upon  the  formation  of  the  cuprous 
iodide. 

These  points  have,  therefore,  been  carefully  investigated  ex- 
perimentally by  Gooch  and  Heath,*  with  the  results  summarized 
below. 

As  to  the  use  of  potassium  iodide  in  effecting  the  precipitation 
of  cuprous  iodide,  it  appears  that  the  excess  present  has  within 
limits  an  influence  upon  the  result;  that  beyond  the  limits  the 
addition  of  potassium  iodide  has  no  appreciable  effect;  and  that 
the  absolute  amount  of  potassium  iodide  required  increases  with 
the  dilution.  An  excess  of  potassium  iodide  ranging  from  0.6 
grm.  to  I  grm.  in  a  volume  of  50  cm.3,  and  from  3  grm.  to  5  grm. 
in  a  volume  of  100  cm.3,  will  precipitate  completely  0.0020  grm. 
of  copper.  In  the  practical  application  of  these  facts  it  must  be 
borne  in  mind  that  it  is  the  excess  of  potassium  iodide  and  not 
the  full  amount  added  which  is  important.  So  it  is  reasonable 
to  fix  upon  2  grm.  as  the  uniform  amount  of  potassium  iodide 
suitable  for  the  precipitation  of  cuprous  iodide  equivalent  to 
0.2  grm.  of  copper  in  a  volume  of  50  cm.3;  and  upon  5  grm.  as 
the  amount  of  potassium  iodide  suitable  in  neutral  solutions 
having  a  volume  of  100  cm.3. 

In  a  study  of  the  effect  of  free  acid  upon  potassium  iodide 
alone,  it  appears  that  no  more  than  2  cm.3  of  concentrated  sul- 
phuric acid  or  hydrochloric  acid  may  be  present  with  2  grm. 
of  potassium  iodide  in  50  cm.3  of  solution  without  liberating  an 
appreciable  amount  of  iodine,  and  the  presence  of  I  cm.3  of  pure 
nitric  acid  causes  error.  The  tendency  to  liberate  iodine  is 
manifestly  less  at  the  higher  dilution,  and  it  appears  that  in  a 
volume  of  100  cm.3  of  solution  containing  5  grm.  of  potassium 
iodide  3  cm.3  of  concentrated  sulphuric  acid,  hydrochloric  acid 
or  nitric  acid  free  from  nitrogen  oxides  may  safely  be  present. 
Acetic  acid  of  50  per  cent  strength  may  apparently  make  up 
half  the  solution  at  either  dilution.  When  either  sulphuric  acid, 
hydrochloric  acid  or  nitric  acid  is  present,  obviously  the  higher 
dilution  is  preferable.  In  the  determination  of  o.ooio  grm.  of 
copper  by  titration  of  the  iodine  set  free  in  a  volume  of  100  cm.3 
in  presence  of  5  grm.  of  potassium  iodide,  it  appears  that  so 
much  as  50  cm.3  of  50  per  cent  acetic  acid,  3  cm.3  of  sulphuric 
*  F.  A.  Gooch  and  F.  H.  Heath,  Am.  Jour.  Sci.,  [4],  xxiv,  67. 


122 


METHODS  IN  CHEMICAL  ANALYSIS 


lodometric  Determination  of  Copper. 


KI. 

Volume. 

Copper 
taken  as 
Cu(NO,)2. 

Acid. 

Copper 
found. 

Error. 

Present. 

Approx- 
imate 

At  begin- 
ning of 

At  end 
of  titra- 

excess. 

titration. 

tion. 

grrn. 

grm. 

grm. 

cm.s 

cm.* 

cm.  s 

grm. 

grm. 

Final  volume  between  no  cm.3  and  120  cm.3. 


HaSO4  cone. 

0.1200 
O.OQOO 
O.O9OO 

5-o 
5-o 
S-o 

4-5 
4-5 
4-5 

2-5 
3-0 
3-5 

100 
100 
IOO 

119 
114 
114 

O.  I2OO 

0.0903 
0.0905 

o.oooo 
+0.0003 
4-0.0005 

HC1  cone. 

0.0900 

5-o 

4-5 

2.O 

IOO 

117 

0.0897 

—0.0003 

0.1200 
O.O9OO 

S-o 
5-o 

4-5 
4-5 

2.0 

3-o 

IOO 

IOO 

119 
114 

0.1195 

o  .  0901 

—0.0005 

+  0.0001 

O.I2OO 

5-0 

4-5 

3-o 

IOO 

119 

O.  I2OO 

o.oooo 

O.I2OO 
0.0900 

S-o 
S-o 

4-5 
4-5 

3-5 
4-o 

IOO 
IOO  ' 

119 

114 

0.1197 

o  .  0903 

—0.0003 
+0.0003 

HNO,  cone. 

0.09.00 
O.IO5O 
0.0900 

S-o 
S-o 
S-o 

4-5 
4-5 
4-5 

i  .0 
i-5 
2-5 

IOO 
IOO 
IOO 

114 
117 
114 

0.0900 
0.1051 
0.0901 

o.oooo 

+  O.OOOI 
+  O.OOOI 

50  per  cent 
HC2H302. 

O.I2OO 
0.0900 
0.1050 

S-o 
S-o 
S-o 

4-5 
4-5 
4-5 

3-0 
S-o 
IO.O 

IOO 
IOO 
IOO 

119 
114 
117 

0.1195 

0.0898 

0.1048 

—0.0005 

—  O.OOO2 
—  O.OOO2 

Final  volume  increased  by  titration  to  132  cm.3  and  150  cm.3;  KI  corre- 
spondingly increased. 


H2SO4  cone. 

0.2218 

0.3231 

7-o 
8.0 

6.0 
6.4 

2 

3 

IOO 
IOO 

135 
150 

0.2214 

0.3226 

—0.0004 

-0.0005 

HC1  cone. 

0.2023 
0.2581 

7-0 
7-8 

6.0 
6.7 

2 

3 

IOO 
IOO 

132 
141 

0.2016 

0.2574 

—  0.0007 
—0.0007 

HNOj  cone, 
purified. 

0.2023 
0.2520 

8.0 

IO.O 

7-o 
8-5 

I 
3 

IOO 
IOO 

132 
148 

0.2017 

0.2512 

—  0.0006 
—0.0008 

HC2H30, 
50  per  cent. 

0.2125 
0.2064 

7-5 
8.0 

8 

IOO 
IOO 

133 
132 

o.  2119 

o.  2058 

—0.0006 
—0.0009 

COPPER;    SILVER;   GOLD  123 

acid,  3  cm.3  of  hydrochloric  acid,  or  3  cm.3  of  nitric  acid  (free 
from  nitrogen  oxides)  may  be  present  without  appreciable  influ- 
ence upon  the  indications  of  the  process. 

The  best  general  procedure  in  determining  by  the  iodometric 
method  amounts  of  copper  not  exceeding  about  0.3  grm.  seems 
to  be  covered  by  the  following  directions:  The  solution  of  the 
cupric  salt,  containing  no  more  than  3  cm.3  of  concentrated 
sulphuric  acid,  hydrochloric  acid  or  nitric  acid  (free  from  nitro- 
gen oxides),  or  25  cm.3  of  50  per  cent  acetic  acid,  is  made  up  tc* 
a  volume  of  100  cm.3,  5  grm.  of  iodate-free  potassium  iodide 
added,  and  the  titration  of  the  free  iodine  made  by  sodium  thio- 
sulphate  in  the  usual  manner  with  the  use  of  the  starch  indicator 
at  the  end.  The  n/io  sodium  thiosulphate  used  in  the  titration 
adds  appreciably  to  the  initial  100  cm.3  of  solution  when  much 
copper  is  estimated,  and  in  case  the  end  reaction  has  not  appeared 
when  so  much  as  25  cm.3  of  the  thiosulphate  has  been  added,. 
2  grm.  to  3  grm.  more  of  potassium  iodide  should  be  added  before 
continuing  the  titration. 

The  error  of  the  process,  properly  conducted,  should  not  exceed' 
a  few  tenths  of  a  milligram  in  terms  of  copper. 

Results  obtained  by  this  procedure  are  given  in  the  table. 

The  Determination  of  Copper  by   Titration  of  the  Precipitated 
Oxalate  with  Potassium  Permanganate. 

Upon  the  well-known  fact  that  copper  oxalate  is  insoluble  in 
water  and  scarcely  attacked  by  moderate  amounts  of  dilute  nitric 
acid,  Bournemann*  has  based  an  approximative  method  for  sepa- 
rating copper  from  cadmium,  by  precipitation  as  the  oxalate  in. 
the  presence  of  nitric  acid,  and  estimating  the  copper,  after  igni- 
tion, by  any  of  the  well-known  gravimetric  methods.  Classen  \ 
describes  a  method  for  the  separation  of  metals  as  oxalates  by- 
adding  to  the  solution  of  the  salt  of  the  metals  a  dilute  solution 
of  potassium  oxalate  [i  :  6]  and  concentrated  acetic  acid  to  80 
per  cent  of  the  total  volume.  Regarding  copper  salts  in  par- 
ticular, Classen  states  that  precipitation  takes  place  only  in 
dilute  solution,  and  then  not  completely. 

According  to  the  experience  of  Peters, J  the  precipitation  of 

*  Chem.  Ztg.,  xxiii,  565. 

t  Ber.  Dtsch.  Chem.  Ges.,  x.  b,  1316. 

t  Charles  A.  Peters,  Am.  Jour.  Sci.,  [4],  x,  359. 


124 


METHODS  IN  CHEMICAL  ANALYSIS 


copper  oxalate  from  solutions  saturated  with  oxalic  acid  and 
containing  at  least  o.oi  grm.  of  copper  in  50  cm.3  of  liquid  may 
be  made  practically  complete,  the  filtrate  in  such  cases  giving 
no  blue  color  with  ammonia,  in  a  test  tube  viewed  lengthwise, 
and  only  a  faint  brown  color  when  the  filtrate  is  neutralized,  made 
acid  with  acetic  acid,  and  tested  with  potassium  ferrocyanide. 
Peters  shows  that  copper  may  be  determined  quantitatively  as 
the  oxalate,  by  precipitation  with  oxalic  acid  and  titration  of 
the  precipitate  by  potassium  permanganate,  and  separated  from 
certain  other  metals  in  the  presence  of  nitric  acid,  by  the  addi- 
tion of  considerable  amounts  of  oxalic  acid,  provided  that  the 
amount  of  copper  present  in  solution  exceeds  a  certain  mini- 
mum value.  It  is  shown  that  when  the  amount  of  copper 
present  falls  below  the  minimum  either  precipitation  does  not 
take  place  or  it  is  incomplete.  It  is  noted  that  the  minimum  is 
variable  with  the  concentration  of  the  precipitant,  oxalic  acid, 
and  to  some  extent  dependent  upon  the  condition  of  the  pre- 
cipitant, the  minimum  being  smaller  when  the  oxalic  acid  is 
added  in  crystalline  form  rather  than  in  solution  to  the  liquid 
containing  the  copper  salt.  Peters'  observations  in  respect  to 
the  effect  of  concentration  and  condition  of  the  oxalic  acid  in 
solution  may  be  summarized  in  the  following  statement: 


Minimum  amount  of 
copper,  taken  as  the 

Amount  of  oxalic  acid  used. 

be  present  in  order 
that  nearly  complete 
precipitation  may 

Crystalline. 

In  solution. 

Volume  of 
liquid. 

take  place. 

grm. 

grm. 

grm. 

cm.» 

O.OIO 

5 

5* 

50 

0.025 

2 

3-5 

SO 

0.040 

I 

50 

0.050 

0-5 

50 

•  Saturated  solution  added  to  the  copper  salt  dissolved  in  the  least  amount  of  water. 

It  is  noted  that  when  a  saturated  solution  of  oxalic  acid,  con- 
taining o.i  grm.  of  oxalic  acid  to  I  cm.3,  is  slowly  added  to  a  drop 
of  the  copper  solution  containing  0.0003  grm-  of  copper,  the 
precipitated  oxalate  first  formed  dissolves  completely  in  a  volume 
of  5  cm.3  of  the  precipitant. 

When  no  added  nitric  acid  is  present  precipitates  formed  in 
the  hot  solutions  at  a  volume  of  50  cm.3  may  be  filtered,  either 


COPPER;    SILVER;    GOLD 


125 


at  once  or  after  cooling,  without  loss;  but  the  condition  of  the 
precipitate  is  better  in  the  presence  of  nitric  acid.  When  the 
nitric  acid  is  added  the  mixture  must  stand  before  filtration  — 
best  over  night.  Results  are  unsatisfactory  when  ammonium 
nitrate  is  present. 

Precipitation  of  According  to  the  procedure  recommended  by  Pe- 
Copper  oxaiate.  tergj  COpper  sulphate  dissolved  in  approximately  50 
cm.3  of  hot  water,  to  which  nitric  acid  is  added  to  the  amount 
of  5  cm.3  when  certain  separations  are  to  be  effected,  is  treated 
in  a  small  beaker  with  crystallized  oxalic  acid,  2  grm.  to  3  grm., 
and  allowed  to  stand  over  night.  The  precipitate  is  filtered  off 
on  asbestos  and  washed  two  or  three  times  with  small  amounts 
of  cold  water  and,  still  in  the  crucible,  is  returned  to  the  beaker, 
5  cm.3  or  10  cm.3  of  dilute  sulphuric  acid  [i  :  i]  are  added,  to- 
gether with  a  convenient  amount  of  water,  and,  after  heating  the 
liquid  to  boiling,  the  oxalic  acid  is  titrated  with  permanganate, 
the  oxaiate  of  copper  dissolving  readily  as  fast  as  the  excess  of 
oxalic  acid  is  removed  by  the  permanganate.  The  precipitate 
may  also  be  dissolved  in  10  cm.3  of  strong  hydrochloric  acid  and 
titrated  at  3O°-5O°  after  the  addition  of  0.5  grm.  of  manganous 
chloride.  Experimental  results  are  given  below. 

Permanganate  Titration  of  Copper  Oxaiate. 


CuO  taken  as 
CuSO4. 

grm. 

Oxalic 

acid. 

grm. 

HNO3 
(sp.  gr.  1.40). 

cm.3 

Volume  at 
precipitation. 

cm.3 

CuO  found, 
grm. 

Error, 
grm.' 

o.  1860 

o-5 

50 

0.1864 

+0.0004* 

0.1860 

0-5 

50 

0.1866 

+O.OOO6 

0.1860 

o-S 

5° 

0.1866 

+0.0006 

0.1860 

I.O 

50 

0.1866 

+0  .  0006 

0.0398 

I.O 

50 

0.0391 

—  0.0007 

0.1860 

o-5 

5-o 

55 

0.1859 

—  O.OOOI 

0.1860 

o-5 

5-o 

55 

0.1860 

o.oooo 

o.  1990 

2.O 

5-o 

55 

o.  1989 

—  O.OOOI 

o.  1990 

3-o 

5-o 

55 

o.  1990 

o.oooo 

*  The  titration  was  made  in  the  hydrochloric  acid  solution  containing  manganous  chloride. 

Solubility  of  The  fact  noted  by  Peters,  that  small  amounts  of 

Copper  oxaiate.  precipitated  copper  oxaiate  may  be  redissolved  in  a 
sufficient  excess  of  the  precipitant,  points  to  an  appreciable  de- 
gree of  solubility  of  the  precipitate  in  the  solution  of  oxalic  acid. 
The  observation  that  very  considerable  amounts  of  copper  oxa- 


126  METHODS  IN  CHEMICAL  ANALYSIS 

late  fail  to  come  down  at  all  until  a  certain  minimum  of  the  copper 
salt  is  present,  although  precipitation  is  nearly  complete  when 
that  minimum  is  reached,  indicates  supersaturation  of  the  so- 
lution by  copper  oxalate;  while  the  capacity  of  the  liquid  for 
supersaturation  is  apparently  limited  to  some  extent  by  increase 
in  concentration  of  the  oxalic  acid.  The  solubility  coefficient  of 
the  copper  oxalate  under  the  conditions  is  made  up,  therefore, 
of  at  least  two  factors,  of  which  one  depends  upon  the  normal 
solubility  in  the  solution  of  oxalic  acid  which  constitutes  the 
medium  of  precipitation,  while  the  other  depends  upon  the  solu- 
bility due  to  supersaturation.  In  order  that  small  amounts  of 
copper  may  be  precipitated,  it  is  necessary  to  eliminate,  or  at 
least  to  limit,  the  capacity  of  the  medium  for  supersaturation; 
and  in  order  that  large  amounts,  as  well  as  small  amounts,  of 
copper  may  be  determined  with  the  highest  degree  of  accuracy, 
it  is  necessary  to  reduce  to  the  lowest  point  the  normal  solubility 
of  the  oxalate  under  the  conditions  of  precipitation. 

Gooch  and  Ward  *  have  studied  the  conditions  under  which 
small  as  well  as  large  amounts  of  copper  may  be  determined  by 
the  oxalate  method.  It  is  to  be  noted,  in. the  first  place,  that  the 
character  of  precipitated  copper  oxalate  depends  upon  the  con- 
ditions of  precipitation.  When  oxalic  acid  is  added  to  a  cold 
concentrated  solution  of  a  salt  of  copper,  the  copper  oxalate 
precipitated  is  of  extreme  fineness  and  tends  to  pass  through  the 
closest  filters.  The  precipitate  formed  in  hot  solution  is,  on  the 
other  hand,  crystalline  and  easily  separated  by  filtration  of  this 
liquid.  The  solubility  of  the  precipitate,  as  well  as  the  ease 
with  which  it  may  be  separated  from  the  liquid,  turns  upon  the 
conditions  of  precipitation  and  treatment.  Throughout  a  series 
of  experiments  the  error  of  the  determination  increases  with  the 
dilution.  That  the  errors  found  in  titration  actually  represent 
approximately  the  losses  in  copper  for  the  smaller  volumes,  is 
shown  by  the  electrolytic  determination  of  copper  in  the  filtrates 
from  the  precipitated  oxalate.  For  a  volume  of  10  cm.3,  the 
average  error  in  the  titration  of  the  oxalate  precipitated,  either 
from  the  solution  of  the  sulphate  or  from  a  solution  of  the  nitrate, 
is  0.0002  grm. ;  for  50  cm.3  it  is  o.ooi  I  grm. ;  for  100  cm.3,  0.0053 
grm.;  for  200  cm.3,  0.0203  grm.  For  similar  concentrations  of 
the  copper  salt  and  of  the  oxalic  acid,  the  deficiency  in  the  copper 
*  F.  A.  Gooch  and  H.  L.  Ward,  Am.  Jour.  Sci.,  [4],  xxvii,  448. 


COPPER;    SILVER;    GOLD 


127 


indicated  by  titration  of  the  precipitated  oxalate  increases  more 
rapidly  than  the  dilution,  a  fact  which  suggests  some  specific 
action  of  water,  —  perhaps  hydration  affecting  the  solubility, 
or  hydrolysis  affecting  the  composition,  of  the  copper  oxalate. 
That  time  and  temperature  are  not  essential  factors  in  the  pre- 
cipitation of  the  oxalate  at  moderate  dilution  from  solutions  of 
the  neutral  salt,  was  shown  by  Peters  in  the  following  experi- 
ments which  indicate  also  that  the  precipitates,  whether  thrown 
down  in  hot  solution  or  in  cold  solution,  possess  after  long 
standing  the  same  degree  of  insolubility. 

Effects  of  Temperature  at  Precipitation  and  Filtration  ;  Filtration  after  Stand- 
ing over  Night. 


Copper 
taken. 

Volume. 

Oxalic 
acid. 

Copper 
found. 

Error. 

Precipitation. 

Filtration. 

grni. 

cm.s 

grm. 

grm. 

grm. 

0.0502 

50 

2.0 

0.0491 

O.OOII 

Hot. 

Cold. 

0.0502 

5° 

2.O 

0.0492 

O.OOIO 

Hot. 

Hot. 

0.0502 

50 

2.O 

o  .  0490 

O.OOI2 

Cold. 

Hot. 

0.0502 

50 

2.O 

0.0491 

O.OOII 

Cold. 

Cold. 

If  any  part  of  the  apparent  loss  of  copper  oxalate  precipitated 
from  solutions  of  oxalic  acid  is  due  to  hydrolysis  of  the  normal 
oxalate,  and  formation  of  a  basic  oxalate  as  the  product  of  hydro- 
lytic  action,  it  should  be  possible  to  obviate  such  apparent  loss 
by  increasing  the  active  acidity  of  the  solution  and  thus  inhibit- 
ing hydrolysis,  provided  the  solubility  of  the  normal  oxalate  is 
not  made  greater  thereby.  Experiment  shows  that  beyond  a 
reasonable  degree  of  concentration  the  results  are  not  affected 
by  the  use  of  oxalic  acid  up  to  the  point  of  saturation  of  the  solu- 
tion, but  that  the  apparent  error  is  actually  diminished  by  the 
presence  of  even  very  small  amounts  of  sulphuric  acid  or  nitric 
acid  in  the  liquid,  while,  within  reasonable  limits,  the  addition  of 
more  acid  produces  no  further  effect.  At  the  higher  dilution  the 
effect  of  the  active  acid  is  marked.  At  a  volume  of  100  cm.3  the 
average  error  of  deficiency  shown  in  absence  of  the  stronger  acids 
is  cut  in  two  by  the  addition  of  o.i  cm.3  to  5  cm.3  of  nitric  acid, 
or  of  0.5  cm.3  to  2  cm.3  of  sulphuric  acid.  At  the  smaller  vol- 
ume of  50  cm.3  the  effect  is  not  so  marked,  but  it  is  still  obvious. 
These  results  favor  strongly  the  hypothesis  that  copper  oxalate 


128  METHODS  IN  CHEMICAL  ANALYSIS 

is  increasingly  subject  to  hydrolysis  as  dilution  increases,  and 
that  the  tendency  to  form  a  basic  salt  may  be  checked  by  the 
presence  of  the  stronger  acids  in  suitable  amounts.  Even  very 
large  amounts  of  nitric  acid  produce  a  surprisingly  small  increase 
in  the  apparent  solubility  of  the  oxalate. 

Losses  due  to  solubility  of  copper  oxalate  may  evidently  be 
kept  at  low  limits  by  restricting  the  volume  of  the  solution  of 
oxalic  acid  in  which  precipitation  takes  place;  but  too  high 
concentration  is  likely  to  introduce  error  due  to  mechanical  in- 
clusion of  oxalic  acid  in  the  precipitate.  The  natural  alternative 
to  a  close  restriction  of  the  volume  of  the  aqueous  solution  is  the 
limitation  of  the  solvent  power  of  a  larger  volume  of  liquid  by 
partially  substituting  for  water  some  other  miscible  liquid  less 
capable  of  dissolving  the  precipitated  oxalate.  In  testing  the 
effect  of  alcohol,  suggested  by  Gibbs,*  it  was  found  that  the 
presence  of  that  medium,  either  with  or  without  nitric  acid,  im- 
proves the  results  obtained  at  similar  dilutions  of  the  oxalic  acid 
solution,  and  to  about  the  same  degree  whether  with  or  without 
nitric  acid.  So  it  would  seem,  if  the  effect  of  nitric  acid  is  to 
prevent  the  formation  of  a  basic  salt,  that  alcohol  not  only 
diminishes  the  normal  solubility,  but  checks  hydrolytic  action 
as  well.  In  a  volume  of  100  cm.3  containing  20  per  cent  of  alco- 
hol the  error  approximates  — o.ooio  grm.;  and  for  a  volume  of 
50  cm.3  containing  50  per  cent  alcohol  the  error  is  reduced  to 
—0.0003.  The  effect  of  nitric  acid  accompanying  the  alcohol 
is  not  marked. 

In  further  experiments  it  was  found  that  the  addition  of  acetic 
acid,  as  proposed  by  Classen,  f  is  even  more  effective  than  the  use 
of  alcohol,  or  of  alcohol  with  nitric  acid,  but  that  when  consid- 
erable amounts  of  copper  are  present  the  precipitates  formed  in 
solutions  containing  acetic  acid  are  apt  to  be  very  finely  divided 
and  consequently  difficult  to  filter.  A  better  condition  of  the 
precipitate  is  obtained,  however,  if,  with  the  acetic  acid,  there  is 
also  present  a  moderate  amount  of  nitric  acid.  It  appears  that 
acetic  acid,  when  present  to  the  amount  of  25  per  cent  of  the 
liquid,  produces  in  volumes  of  100  cm.3  about  the  same  effect  as 
alcohol,  and,  when  present  to  the  amount  of  50  per  cent,  diminishes 
still  further  the  solvent  power  of  the  medium  for  the  oxalate. 

*  Am.  Jour.  Sci.,  [2],  xliv,  214. 

f  Ber.  Dtsch.  chem.  Ges.,  x,  b,  1316. 


COPPER;   SILVER;    GOLD  129 

The  additional  presence  of  nitric  acid  to  10  per  cent  of  the  entire 
volume  does  not  materially  affect  the  solubility.  Sulphuric  acid 
to  10  per  cent  of  the  volume  of  the  liquid  is  without  apparent 
effect  upon  the  solubility  of  copper  oxalate,  provided  the  oxalic 
acid  is  also  present  in  the  proportion  of  4  grm.  to  100  cm.3  of  the 
liquid.  Treatment  by  oxalic  acid  in  a  medium  consisting  of  acetic 
acid  of  half  strength,  with  or  without  nitric  to  the  extent  of  10  per 
cent  by  volume,  is  plainly  the  best  of  the  procedures  studied  for 
the  complete  precipitation  of  copper  oxalate  in  ideal  condition. 
Prevention  of  Gooch  and  Ward*  have  made  use  of  various  means 
Supersaturation.  m  ^e  effort  to  break  up  supersaturation  of  the  pre- 
cipitating medium  when  only  small  amounts  of  copper  oxalate 
are  present.  The  supersaturated  solution  was  frozen,  and  the 
mass  melted,  following  procedure  which  has  been  found  to  be 
successful  in  hastening  the  deposition  of  small  amounts  of  am- 
monium magnesium  arsenate;f  the  supersaturated  solution  was 
evaporated  to  dryness,  and  the  residue  extracted  with  water; 
alcohol  was  added  to  the  solution  of  the  copper  salt  before  at- 
tempting precipitation  by  oxalic  acid ;  acetic  acid  of  50  per  cent 
strength  was  used  as  the  medium  in  which  precipitation  was 
attempted  by  oxalic  acid.  From  the  experimental  results  it 
appears  that  by  precipitating  at  a  volume  of  50  cm.3,  freezing, 
melting,  and  boiling,  the  condition  of  supersaturation  may  be 
broken  up,  the  oxalate  obtained  being  soluble  in  the  proportion 
of  about  o.oon  grm.  to  50  cm.3  of  liquid;  that  by  precipitation 
at  a  volume  of  50  cm.3,  evaporation  to  dryness,  and  extraction 
with  the  same  volume  of  water,  the  copper  may  be  recovered  to 
an  amount  within  about  0.0004  grm-  °f  that  taken;  that  treat- 
ment by  oxalic  acid  in  50  per  cent  alcohol  fails  to  precipitate  about 
0.0020  grm.  of  copper  from  amounts  less  than  0.0200  grm.,  while 
for  amounts  exceeding  that  limit  the  copper  is  nearly  all  re- 
covered; and  that  in  volumes  of  50  cm.3  or  100  cm.3,  consisting 
of  50  per  cent  acetic  acid,  the  copper  oxalate  is  thrown  down 
completely,  the  presence  of  nitric  acid  to  the  extent  of  10  per 
cent  making  the  nitration  more  effective  without  influencing  the 
solubility,  while  even  at  a  volume  of  150  cm.3  the  precipitation 
is  complete  provided  the  acetic  acid  makes  up  two-thirds  of  the 
volume.  Acetic  acid  appears,  therefore,  to  be  most  effective  in 

*  Loc.  cit. 

t  Gooch  and  Phelps,  see  page  290. 


130 


METHODS  IN  CHEMICAL  ANALYSIS 


Presence  of 
Acetic  Acid. 


breaking  up  the  condition  of  supersaturation  as  well  as  in  dimin- 
ishing the  degree  of  normal  solubility  in  the  medium  of  pre- 
cipitation. 

Precipitation  in  According  to  the  procedure  recommended  by 
Gooch  and  Ward,  for  small  amounts  as  well  as  the 
larger  amounts  of  copper,  oxalic  acid,  2  grm.  or 

4  grm.,  is  added  to  the  copper  salt  dissolved  in  50  cm.3  or  100  cm.3, 
respectively,  of  the  50  per  cent  solution  of  acetic  acid  containing 

5  per  cent  to  10  per  cent  of  nitric  acid  to  induce  a  favorable  con- 
dition for  crystallization.     After  standing  over  night  in  contact 
with  the  solution,  the  precipitate  is  collected  upon  asbestos  in  a 
perforated  crucible  and  washed  carefully  with  small  amounts  of 
water.     The  crucible  with  its  contents  is  placed  in  a  beaker  and 
covered  with  about  200  cm.3  of  hot  water  containing  25  cm.3 
of  dilute  sulphuric  acid  [i  13],  and  the  solution  is  titrated  with 
n/io  potassium  permanganate.     Results  of  this  procedure  are 
given  in  the  table. 

Filtration  of  Copper  Oxalate  Precipitated  in  Solutions  Containing  50  per  cent 
Acetic  Acid  and  5  to  10  per  cent  Nitric  Acid. 


Copper 
taken. 

grm. 

Total 
volume. 

cm.3 

Oxalic 
acid. 

grm. 

Acetic 
acid. 

cm.3 

Nitric 
acid. 

cm.8 

Copper  found, 
grm. 

Error, 
grm. 

Volume  50  cm.3. 


O.OOIO 

50 

2 

25 

5 

O.OOIO 

o.oooo 

O.OO2O 

SO 

2 

25 

5 

O.OO2I 

+O.OOOI 

0.0031 

SO 

2 

25 

5 

0.0033 

+  O.OOO2 

O.OO4I 

50 

2 

25 

5 

o  .  0042 

+0.0001 

O.OO5I 

50 

2 

25 

5 

0.0049 

—  0.0002 

O.OIO2 

50 

2 

25 

5 

0.0103 

+  O.OOOI 

o  .  0204 

50 

2 

25 

5 

o  .  0204 

o.oooo 

0.05II 

50 

2 

25 

5 

0.0512 

+O  .  OOOI 

Volume  100  cm.3. 

0.0031 

IOO 

4 

50 

5 

0.0031 

0.0000 

0.0041 

IOO 

4 

50 

5 

0.0041 

0.0000 

0.0051 

IOO 

4 

So 

5 

0.0051 

o.oooo 

0.0102 

IOO 

4 

50 

5 

0.0103 

+O.OOOI 

o  .  0204 

IOO 

4 

So 

5 

0.0196 

-0.0008* 

0.05II 

IOO 

4 

SO 

5 

0.0510 

—  o.oooi 

0.05II 

no 

4 

50 

10 

o  .  0506 

—0.0005 

0.05II 

IOO 

4 

50 

10 

0.0510 

—o.oooi 

0.1530 

IOO 

4 

50 

10 

0.1529 

—o.oooi 

0.1530 

IOO 

4 

SO 

10 

0.1530 

o.oooo 

Filtration  imperfect. 


COPPER;   SILVER;   GOLD 


Separations  by  Peters  has  shown  *  that  copper  exceeding  the  mini- 
Peters'  Proce-  mum  amount  of  o.oi  grm.  in  50  cm.3  (that  is,  in 

dure  from  Cad-  .  .  .     .  , 

mium,  Arsenic,  amount  sufficient  to  be  precipitated  |  according  to 
iron, Tin, zinc,  ^is  procedure),  may  be  successfully  separated  from 
cadmium,  arsenic,  and  iron  taken  as  ferric  nitrate.  When  iron 
is  present  as  the  sulphate  the  results  are  low.  The  separation 
from  tin  in  the  stannous  form  is  fairly  good  for  amounts  of  that 
element  not  exceeding  o.i  grm.  In  presence  of  the  stannic  salt 
the  losses  are  considerable.  Attempted  separations  from  bis- 
muth and  antimony  were  unsuccessful,  and  the  separation  from 
zinc  unsatisfactory  on  account  of  the  tendency  of  zinc  oxalate  to 
fall  with  the  copper  oxalate.  Experimental  results  are  given  in 
the  table. 

Permanganate  Titration  of  Copper  Oxalate  in  Separations. 


CuO 

taken  as 
CuSO4. 

grm. 

Element  from 
which  copper 
was  separated. 

grm. 

Oxalic 

acid. 

grm. 

HNO3 

(sp.  gr.  1.40)  . 

cm.8 

Volume  at 
precipita- 
tion. 

cm.3 

CuO 

found. 

grm. 

Error, 
grm. 

Cadmium. 


CdO  taken 

/ 

as  CdSO4. 

o.  1990 

O.IO 

2.0 

5-o 

60 

0.1983 

—0.0007 

0.1990 

O.2O 

2.0 

5-o 

65 

0.1987 

—0.0003 

o.  1990 

0.30 

2.O 

5-o 

70 

0.1987 

—0.0003 

o.  1990 

0.40 

2.O 

5-o 

75 

0.1994 

+0.0004 

0.1990 

0.50 

2.O 

S-o 

80 

0.1996 

+0.0006 

Arsenic. 


As2O3  taken 

as  Na3AsO3. 

,. 

0.1990 

O.IO 

2.0 

55 

0.1991 

+0.0001 

0.1990 

O.2O 

2.0 

60 

0.1987 

—0.0003 

o.  1990 

0.50 

2.O 

75 

0.1986 

—0.0004 

0.1990 

O.2O 

2.O 

5-o 

60 

0.1994 

+o  .  0004 

0.1990 

O.2O 

2.O 

5-o 

75 

0.1992 

+O.OOO2 

0.1990 

0.60 

2.O 

5-o 

85 

0.1995 

+0.0005 

AssO,  taken 

as  H2KAs04. 

0.1990 

O.IO 

2.O 

.  .  . 

60 

0.1985 

—0.0005 

0.1990 

O.2O 

2.O 

70 

0.1990 

0.0000 

0.1990 

O.IO 

2.0 

5-o 

65 

0.1990 

0.0000 

0.1990 

O.2O 

2.0 

S-o 

75 

0.1992 

+O.OOO2 

0.1990 

0.30 

.       2.0 

5-o 

85 

0.1985 

—0.0005 

o  .  2030 

0.30 

3-o 

S-o 

85 

0.2026 

—0.0004 

See  page  125. 


t  See  page  124. 


I32 


METHODS  IN  CHEMICAL  ANALYSIS 


Permanganate  Titration  of  Copper  Oxalate  in  Separations. 


CuO 

taken  as 
CuSO4. 

Element  from 
which  copper 
was  separated. 

Oxalic 
acid. 

HNO3 
(sp.  gr.  1.40). 

Volume  at 
precipita- 
tion. 

CuO 
found. 

Error. 

grm. 

grm. 

grm. 

cm.8 

cm.8 

grm. 

grm. 

Iron. 


Fe2O3  taken 
as  Fe(NO3)3. 

0.1990 

0.136 

2.O 

50 

60 

0.1987 

-0.0003 

0.1990 

o.  272 

2.0 

5-o 

60 

0.1983 

—0.0007 

0.1990 

0.364 

2.0 

5-o 

60 

0.1988 

—  O.OO02 

o.  1990 

0-544 

2.O 

5-o 

65 

0.1971 

—0.0019 

Tin. 


Sn  taken  as 
SnCU-t-HCl. 

Cu  found. 

0.1990 

0  .  0468 

2.O 

5-o 

55 

0.1979 

—  o.oon 

0.1990 

0.0936 

2.0 

S-o 

60 

o  .  2004 

+0.0014 

0.1990 

0.0936 

2.O 

S-o 

60 

o.  1992 

-f-O  .  OOO2 

0.1990 

0.0936 

2.O 

5-0 

60 

0.1995 

+o  .  0005 

Sn  taken  as 

SnCl4. 

0.1990 

O.  IO 

2.O 

5-o 

55 

0.1979 

—  o.oon 

0.1990 

O.IO 

2.O 

55 

0.1959 

—0.0031 

0.1990 

O.2O 

2.O 

5.0 

55 

0.1974 

—  0.0016 

0.1990 

0.50 

2.O 

5-o 

60 

0.1955 

-0.0035 

Zinc. 


ZnO  taken 

as  ZnSO4. 

0.1990 

0.028 

2.0 

5-o 

60 

o.  2007 

+0.0017 

0.1990 

0.057 

2.O 

5-o 

65 

o  .  2008 

+0.0018 

0.1990 

0.057 

2.O 

5-o 

65 

o  .  2008 

+0.0018 

0.1990 

0.085 

2.O 

5-0 

70 

0.2035 

+0.0045 

Separations  by  Ward*  has  studied  the  effect  of  evaporation  to 
tho  Method  of  dryness,  to  break  up  supersaturation,  in  separations 
Desiccation.  Qj-  COpper  frOm  cadmium,  arsenic,  and  iron.  In  this 
process  the  boiling  solution  is  treated  with  oxalic  acid  and 
nitric  acid.  The  liquid  and  precipitate  are  evaporated  to  dry- 
ness,  and  the  residue  is  extracted  with  cold  water  containing 
not  too  much  nitric  acid  and,  in  separations  from  iron,  more 
oxalic  acid.  The  residual  oxalate  is  filtered  .off  on  asbestos  in  the 
perforated  crucible,  washed  with  small  amounts  of  water,  treated 
*  H.  L.  Ward,  Am.  Jour.  Sci.,  [4],  xxxiii,  423. 


COPPER;   SILVER;   GOLD 


133 


with  boiling  water  containing  sulphuric  acid,  and  titrated  with 
permanganate.  The  details  of  the  preferred  treatment  and  the 
results  are  given  in  the  accompanying  tables. 


Copper  from  Cadmium  by  Desiccation  Process. 


Copper 
present. 

Cad- 
mium 
present. 

Nitric 
acid  at 
precipi- 
tation. 

Volume 
at  pre- 
cipita- 
tion. 

Oxalic 
acid. 

Liquid 
used  in 
extrac- 
tion. 

Nitric 
acid  in 
extrac- 
tion. 

Copper 
found. 

Error. 

grm. 

grm. 

cm.* 

cm.3 

gnu  • 

cm.* 

cm.* 

grm. 

grm. 

0.0051 

O.  IO 

5 

5° 

4 

50 

5 

0.0048 

—  0.0003 

0.0514 

O.OI 

5 

SO 

4 

50 

5 

0.0507 

—0.0007 

0.0514 

0.06 

5 

50 

4 

SO 

5 

o  .  0506 

—  0.0008 

o  .  0504 

O.  IO 

5 

SO 

4 

SO 

5 

0.0502 

—  O.OOO2 

0.0514 

O.  IO 

5 

SO 

4 

50 

5 

0  .  0508 

—  o.ooo5 

0.0514 

O.2O 

5 

50 

4 

50 

5 

O  .  0508 

—0.0006 

0.0516 

0.30 

5 

50 

4 

50 

5 

0.0507 

—0.0009 

0.1542 

O.2O 

5 

SO 

4 

50 

5 

0.1537 

—0.0005 

Copper  from  Arsenic  by  Desiccation  Process. 


Copper 
present. 

Arsenic  as 
arsenate. 

Volume 
at  precip- 
itation. 

Oxahc 
acid. 

Liquid 
used  in 
extraction. 

Nitric 
acid  in 
extraction. 

Copper 

found. 

Error. 

grm. 

grm. 

cm.* 

grm. 

cm.* 

cm.* 

grm. 

grm. 

0.0051 

O.  IO 

50 

50 

2 

0.0047 

—0.0004 

0  .  0504 

0.05 

50 

SO 

2 

0.0499 

—0.0005 

0.0504 

0.05 

50 

50 

2 

0.0501 

-0.0003 

0.0504 

O.  IO 

50 

50 

2 

o  .  0503 

—  o.oooi 

0.0504 

O.  IO 

50 

50 

2 

0.0497 

—0.0007 

o  .  0504 

O.2O 

50 

50 

2 

o  .  0498 

—0.0006 

0-1533 

O.  2O 

50 

50 

2 

0.1528 

—0.0005 

Copper  from  Iron  by  Desiccation  Process. 


Volume 

Oxalic 

Liquid 

Nitric 

Oxalic 

'Copper 
present. 

Iron 
present. 

at  pre- 
cipita- 

acid in 
precipi- 

used in 
extrac- 

acid in 
extrac- 

acid add- 
2d  in  ex- 

Copper 
found. 

Error.  ' 

tion. 

tation. 

tion. 

tion. 

traction 

grm. 

grm. 

cm.* 

grm. 

cm.* 

cm.* 

grm. 

grm. 

grm. 

0.0504 

0.0393 

50 

50 

2 

2 

O  .  0500 

—  0.0004 

0.0504 

0.0393 

50 

50 

2 

2 

0.0501 

—  0.0003 

0.0504 

0.0393 

50 

50 

2 

2 

0.0499 

—  0.0005 

0.0504 

0.0786 

50 

50 

2 

2 

0.0499 

-0.0005 

0.0511 

O.  IOOO* 

50 

50 

2 

3 

o  .  0506 

—  0.0005 

0-1533 

O.  IOOO* 

50 

50 

2 

3 

0.1527 

—O.OOO6 

•  More  than  o.i  grm.  of  iron  apparently  occasions  greater  solubility  of  the  copper  oxalate. 


134 


METHODS  IN  CHEMICAL  ANALYSIS 


Separations  in 
Presence  of 
Acetic  Acid. 


WARD*  investigated  also  the  application  of  the 
process  described  above  for  precipitating  copper  oxa- 
late  in  presence  of  50  per  cent  acetic  acid,  and  of 
nitric  acid  up  to  10  per  cent  to  favor  crystallization,  to  separa- 
tions of  copper  from  cadmium,  arsenic,  iron  and  zinc.  Details 
and  results  are  given  in  the  tables. 

*, 

Copper  from  Cadmium  in  50  per  cent  Acetic  Acid. 


Copper 
present. 

Cadmium 
present. 

Volume 
at  precipi- 
tation. 

Oxalic 
acid. 

Acetic 
acid. 

Nitric 
acid. 

Copper 
found. 

Error. 

grin. 

grin* 

cm.3 

grm. 

cm.3 

cm.3 

grin. 

grm. 

0.0051 

O.2O 

IOO 

4 

50 

5 

0.0049 

—  O.OOO2 

0.0051 

0.30 

IOO 

4 

50 

5 

0.0053 

+  O.OOO2 

0.0255 

O.  2O 

IOO 

4 

50 

10 

0.0257 

+O.OOO2 

0.0510 

0.20 

IOO 

4 

50 

10 

0.0512 

+0  .  0002 

0.0511 

0.30 

IOO 

4 

5° 

IO 

0.0520 

+0.0009 

0.1533 

0.30 

IOO 

4 

50 

IO 

0.1556 

+0.0023 

0.1629 

0.30 

IOO 

4 

50* 

IO 

o.  1636 

+O.OOO7 

*  The  acetic  acid  was  added  after  precipitation.    This  apparently  makes  a  sharper  separation 
of  the  cadmium  from  large  amounts  of  copper. 

Copper  from  Arsenic  in  50  per  cent  Acetic  Acid. 


Copper 

present. 

Arsenic 
present  as 
arsenate. 

Volume 
at  precipi- 
tation. 

Oxalic 
acid. 

Acetic 
acid. 

Nitric 
acid. 

Copper 
found. 

Error. 

grm. 

grm. 

cm.3 

grm. 

cm.3 

cm.3 

grm. 

grm. 

0.0051 

0.30 

IOO 

4 

50 

0.4 

0.0054 

+0.0003 

0.0511 

0.30 

IOO 

4 

50 

0.4 

0  .  0505 

—  0.0006 

0.1530 

0.20 

IOO 

4 

50 

10 

0.1530 

o.oooo 

O.I53S 

0.30 

IOO 

5 

50 

IO 

0.1530 

—0.0005 

Copper  from  Iron  in  50  per  cent  Acetic  Acid. 


Copper 

present. 

Iron 
present. 

Volume 
at  precipi- 
tation. 

Oxalic 
acid. 

Acetic 
acid. 

Nitric 
acid. 

Copper 
found. 

Error. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

0.0510 

0.188 

IOO 

4 

50 

0.0514 

+0.0004 

0.0510 

0.188 

IOO 

4 

50 

2 

0.0511 

+  O.OOOI 

0.0510 

0.188 

IOO 

4 

50 

5 

o  .  0499 

—  O.OOII 

0.0511 

O.  IOO 

IOO 

4 

SO 

IO 

o  .  0489 

—  O.OO22 

0.0510 

0.188 

IOO 

4 

50 

10 

0.0487 

—  0.0023. 

Much  free  nitric  acid  obviously  occasions  loss  of  copper  oxa- 
late.     Ward  t  advocates  as    the    best    procedure   in   separating 

*  Loc.  cit. 


COPPER;    SILVER;    GOLD 


135 


copper  from  iron  the  nearly  complete  precipitation  of  the  copper 
by  oxalic  acid  added  to  the  boiling  solution  devoid  of  other  free 
acid,  with  the  subsequent  addition  of  acetic  acid  in  amount  equal 
to  twice  the  volume  of  the  solution,  to  complete  the  precipita- 
tion, after  cooling.  In  this  manner  the  results  of  the  following 
table  were  obtained. 

Copper  from  Iron  by  Precipitation  in  Water  Solution  with  Subsequent  Addition 

of  Acetic  Acid. 


Copper 
present. 

grm. 

Iron 
present. 

grm. 

Volume  at 
precipita- 
tion. 

cm.3 

Oxalic 
acid. 

grm. 

Acetic 
acid. 

cm.1 

Copper  found, 
grm. 

Error, 
grm. 

O.OO5I 

0.31 

50 

6 

IOO 

0.0049 

—  O.OOO2 

0.0051 

0-45 

50 

6 

IOO 

0.0046 

—  0.0005 

0.0543 

0-15 

50 

6 

IOO 

0.0544 

+0.0001 

0.0543 

0.21 

50 

6 

IOO 

0.0542 

—  o.oooi 

0.0543 

0.31 

50 

6 

IOO 

0.0546 

+0.0003 

0.0543 

0.45 

IOO 

12 

2OO 

0.0538 

—0.0005 

0.1629 

0.45 

50 

6 

IOO 

o.  1649 

+O.OO2O 

0.1629 

o.4S 

IOO 

12 

2OO 

o.  1629 

o.oooo 

When  copper  and  iron  are  -found  together  in  acid  solution,  the 
free  acid  should  either  be  removed  by  evaporation  or  neutralized 
by  potassium  hydroxide  and  the  solution  made  faintly  acid  with 
acetic  acid  before  precipitating  the  copper  as  oxalate.  Ward  has 
shown  that  the  presence  of  moderate  amounts  of  potassium  sul- 
phate, nitrate  or  chloride  does  not  affect  appreciably  the  ana- 
lytical results.  The  presence  of  ammonium  salts,  must,  however, 
be  avoided,  on  account  of  the  tendency  of  copper  oxalate  to 
form  a  soluble  double  ammonium  oxalate. 

Determination  of  The  oxalate  of  lead,  though  fairly  soluble  in  nitric 
Copper  Assod-  acid,  falls  with  copper  oxalate  from  the  nitric  acid 
solution,  and  so  the  separation  of  those  elements  by 
precipitation  of  copper  oxalate  from  the  acid  solution  is  not 
feasible.  Ward*  has  shown,  however,  that  lead  may  be  first 
precipitated  completely  as  lead  sulphate,  and  that  then,  either 
with  or  without  previous  removal  of  the  precipitated  sulphate, 
the  copper  may  be  determined  by  precipitation  and  titration  of 
the  oxalate.  To  the  solution  of  lead  and  copper  in  the  form  of 
nitrates  are  added  an  equal  volume  of  acetic  acid  and  3  cm.3- 
5  cm.3  of  sulphuric  acid.  The  precipitated  sulphate  may  be 

*  Loc.  cit. 


METHODS  IN  CHEMICAL  ANALYSIS 


filtered  off  and  weighed  and  the  copper  estimated  in  the  filtrate 
by  precipitation  and  titration  as  already  described ;  or,  if  a  deter- 
mination of  copper  only  is  desired,  the  precipitation  of  the  copper 
oxalate  may  be  effected  without  removing  the  lead  sulphate,  sul- 
phate and  oxalate  being  filtered  off  together,  treated  with  boiling 
dilute  sulphuric  acid  and  titrated  to  color  with  permanganate. 
Results  of  each  procedure  are  given  below. 

Copper  and  Lead. 


Copper 
present. 

grm. 

Lead 

present. 

grm. 

Sul- 
phuric 
acid. 

cm.' 

Acetic 
acid. 

cm.* 

Volume 
cm.3 

Oxalic 

acid. 

grm. 

Copper 
found. 

grm. 

Error, 
grm. 

Lead 
found. 

grm. 

Error, 
grm. 

Titration  of  copper  oxalate  precipitated  after  separation  of  lead  sulphate. 

0.0511 
0.0511 
0.0511 

o  .  0500 

O.  IOOO 
O.  IOOO 

3 
5 
5 

O  O  O 
to  »o  to 

no 

IOO 
IOO 

2 
2 

2 

0.0513 
0.0508 
0.0508 

+O.OOO2 
-0.0003 
-0.0003 

0.0499 

o  .  0996 

0.0997 

—  O.OOOI 
—  0.0004 
—  0.0003 

Titration  of  copper  oxalate  without  separation  of  lead  sulphate. 

0.0051 
0.0511 
0.0511 
0.0543 
0.0511 
0.0511 

O.  IO22 
0.1086 
0.1533 
0.1533 

0.30 

0.10 

0.25 
0.30 
0.30 
0.40 

0.30 
0.25 

O.2O 
O.2O 

5 
5 
5 
10 

3 
3 

10 

5 
5 
5 

50 
50 
50 
50 
50 
50 
50 
50 
5° 
50 

IOO 
IOO 
IOO 
IOO 
IOO 
IOO 
IOO 
IOO 
IOO 
IOO 

4 

2 
2 

4 

2 

4 

2 

4 

2 
2 

0.0052 
o  .  0508 

O.O5II 
0-0537 
0.0509 

o  .  0508 
0.1018 
o.  1081 
0.1527 
0.1530 

+  O.OOOI 

—0.0003 
o  .  oooo 
—0.0006 

—  O.OOO2 

—0.0003 
—0.0004 
—0.0005 
—  0.0006 
—0.0003 

SILVER. 
The  Gravimetric  Determination  of  Silver  as  the  Chromate. 

The  precipitation  of  silver  chromate  from  the  solution  of 
a  soluble  chromate  made  faintly  acid  with  acetic  acid  may  be 
carried  to  completion  by  the  addition  of  silver  nitrate  in  con- 
siderable excess.  The  exact  determination  of  the  chromium  of  a 
soluble  chromate  or  dichromate  may  therefore  be  effected  by  treat- 
ing with  silver  nitrate  the  solution  of  either  salt,  adding  ammonia 
to  alkalinity  and  then  acetic  acid  to  faint  acidity,  transferring 
the  precipitate  and  washing  it  in  the  filtering  crucible  with  a 
dilute  solution  of  silver  nitrate  until  foreign  material  other  than 
that  reagent  has  been  removed,  finishing  the  washing  with 
a  small  amount  of  water  applied  judiciously  in  portions,  and 


COPPER;   SILVER;    GOLD 


137 


weighing  the  dried  or  gently  ignited  residue  of  silver  chromate.* 
The  success  of  this  process  turns  upon  keeping  the  chromium  at 
the  moment  of  precipitation  essentially  in  the  form  of  chromate 
rather  than  dichromate,  and  in  taking  care  that  an  excess  of 
silver  nitrate  shall  be  present  nearly  to  the  end  of  the  washing. 
Gooch  and  Bosworthf  have  investigated  the  conditions  under 
which,  in  reversal  of  the  process  just  described,  silver  may  be 
precipitated  completely  as  the  chromate,  and  find  that  from 
solutions  of  silver  nitrate  alone  or  containing  free  nitric  acid, 
potassium  chromate  precipitates  silver  chromate  completely, 
provided  enough  potassium  chromate  is  present  to  take  up  the 
nitric  acid  with  formation  of  potassium  nitrate  or  dichromate, 
as  well  as  to  form  the  silver  salt.  The  precipitate,  filtered  at 
once  or  brought  to  better  crystalline  condition  by  dissolving 
it  in  ammonia  and  boiling  the  solution  to  small  volume,  may  be 
transferred  to  the  asbestos  filter  by  dilute  potassium  chromate 

Silver  Weighed  as  Silver  Chromate. 


Silver  taken  as 
AgNO3. 

K2CrO4  used. 

HNO3. 

Ag2CrO4 
weighed. 

Silver 
found. 

Error  in 
terms  of 
silver. 

Volume 

Volume 

of 

Weight. 

of 

Weight. 

Volume. 

Weight. 

solution. 

solution. 

cm.3 

grm. 

cm.3 

grm. 

cm.3 

grm. 

grm. 

grm. 

grm. 

Precipitation  by  K2CrO4. 


is 

0.1652 

50 

0.3 

0.2536 

o.  1649 

—  0.0003 

IO 

O.  IIOI 

50 

0.3 

0.1693 

O.IIOI 

O.OOOO 

25 

0.1437 

50 

0.3 

O.  22OO 

0.1436 

—  o.oooi 

25 

0.1437 

50 

0.3 

O.22IO 

0.1437 

0.0000 

Precipitation  by  K2CrO4  in  presence  of  HNO3. 


25 

0.1355 

50 

o-9 

IO 

0.182 

o.  2091 

0.1360 

+0.0005 

25 

0.1355 

50 

0.9 

IO 

0.182 

0.2081 

0.1353 

—  O.OOO2 

25 

0.1358 

50 

0.9 

IO 

0.182 

0.2090 

0.1360 

+O.OOO2 

25 

0.1355 

50 

0.9 

10 

0.182 

0.2075 

0.1349 

—0.0006 

25 

0.1355 

50 

0.9 

10 

0.182 

0.2090 

0.1360 

+0.0005 

Precipitation  by  K^CrOi  in  presence  of  HNOs,  treatment  with  NH^OH, 
and  boiling  to  a  volume  of  10  to  15  cm.3. 


25 

0.1348 

50 

0.6 

IO 

0.063 

o.  2076 

0.1350 

+O.OOO2 

25 

0.1348 

50 

0.6 

10 

0.063 

0.2068 

0.1344 

—0.0004 

25 

0.1348 

50 

0.6 

10 

0.063 

0.2072 

0.1347 

—  O.OOOI 

25 

0.1348 

50 

0.6 

10 

0.063 

0.2074 

0.1348 

o.oooo 

25 

0.1348 

50 

0.6 

IO 

0.063 

0.2070 

o  .  1346 

—  O.OOO2 

*  See  page  406. 

t  F.  A.  Gooch  and  Rowland  S.  Bosworth,  Am.  Jour.  Sci.,  [4],  xxvii,  241. 


138 


METHODS  IN  CHEMICAL  ANALYSIS 


and  washed  by  small  portions  of  water  without  appreciable  loss. 
The  weight  of  the  residue  of  silver  chromate,  dried  at  gentle  heat, 
may  be  taken  as  a  measure  of  the  silver  originally  present. 

The  Electrolytic  Determination  of  Silver. 

Silver  may  be  determined  by  deposition  of  the  metal  upon  the 
rotary  cathode  *  from  an  ammoniacal  cyanide  solution  in  pres- 
ence of  a  few  grams  of  ammonium  sulphate,  the  method  of  ma- 
nipulation being  precisely  similar  to  that  employed  in  the  depo- 
sition of  copper  as  previously  described.! 

The  following  table  records  determinations  made  in  this  man- 
ner I  with  a  solution  of  silver  nitrate  standardized  as  the  chloride. 

Deposition  of  Silver  from  Ammoniacal  Cyanide  Solution. 


Silver  taken, 
gnu. 

Silver  found, 
grin. 

Error, 
grm. 

Current, 
amp. 

N.  D.100. 

Time, 
min. 

0.0968 

o  .  0966 

—  O.OOO2 

1.8 

6 

IS 

0.0968 

0.0967 

—  O.OOOI 

1.9 

6-3 

15 

0.0968 

o  .  0965 

-0.0003 

1.8 

6 

15 

0.0968 

0.0969 

+O.OOOI 

2 

6.7 

IO 

0.0968 

0.0965 

—0.0003 

3 

10 

8 

0.1898 

o.  1901 

+0.0003 

2.5 

8-3 

10 

0.1898 

0.1898 

o.oooo 

2.5 

8-3 

10 

0.1898 

o  .  1900 

+0.0002 

3 

10 

IO 

0.1898 

0.1893 

—  0.0005 

2-5 

8-3 

IO 

Gooch  and  Feiser  §  have  determined  silver  by  depositing  it 
from  an  ammoniacal  solution  of  the  oxalate,  as  was  done  by 
Gooch  and  Read||  in  the  preparation  of  silver-plated  electrodes, 
in  order  to  avoid  all  possible  contamination  of  the  silver  deposit 
by  nonvolatile  material.  In  testing  this  process,  measured 
amounts  of  the  silver  nitrate  solution  (25  cm.3  or  50  cm.3)  were 
drawn  from  a  burette  into  a  small  beaker  and  treated  with 
ammonium  oxalate  to  complete  precipitation.  The  silver  oxa- 
late was  dissolved  in  a  slight  excess  of  ammonia,  and  this  solution, 
diluted  to  100  cm.3,  was  electrolyzed  with  the  rotary  cathode  and 
a  current  of  0.25  amp.-i.5  amp.  and  4-7  volts.  The  cathode 

*  See  Fig.  13. 

t  See  page  116. 

t  Gooch  and  Medway,  Am.  Jour.  Sci.,  [4],  xv,  320. 

§  F.  A.  Gooch  and  J.  P.  Feiser,  Am.  Jour.  Sci.,  [4],  xxxi,  109. 

II  F.  A.  Gooch  and  H.  L.  Read,  Am.  Jour.  Sci.,  [4],  xxviii,  544. 


COPPER;    SILVER;    GOLD 


139 


with  the  deposited  silver  was  dried  cautiously  over  a  low  flame 
and  thereafter  ignited  to  incipient  redness.  The  details  of 
individual  experiments  are  given  in  the  table. 

Electrolysis  of  Silver  Nitrate  Dissolved  in  Ammonium  Oxalate  and  Ammonia* 


Current. 

Agin 
AgNO3 
taken. 

Ag  found. 

Error. 

Time. 

Revolu- 
tions per 

Amp. 

N?K. 

Volt. 

mm. 

grm. 

grm. 

grm. 

min. 

A  crucible  used  as  cathode. 


0.2687 

0.2685 

—  0.0002* 

i  -5-i 

5-3-3 

6-7 

25 

500 

0.2687 

0.2687 

0.0000* 

i  -5-i 

5-3-3 

6-7 

30 

500 

0.2687 

0.2684 

—  O.OOO3* 

i-o-5 

3-3-1-7 

6-7 

30 

450 

0.2687 

0.2685 

—  O.OOO2* 

1-0.5 

3-3-1-7 

4-6 

30 

450 

0.3183 

0.3181 

—  O.OOO2* 

i-5-i 

5-3-3 

4-6 

20 

450 

0.3183 

0.3178 

—  0.0005t 

i  -5-i 

5-3-3 

4-6 

10 

450 

Gauze  disks  used  as  cathode. 


0.3183 

0.3179 

—0.0004* 

1-0.5 

0.5-0.25 

4-6 

20 

400 

0.3183 

0.3182 

—  O.OOOI* 

1-0.25 

0.5-0.10 

6-8 

25 

400 

0.3183 

0.3181 

—  O.OOO2* 

1-0.25 

0.5-0.10 

5-8 

25 

400 

0.3183 

0.3180 

—  0.0003* 

0.75-0.25 

0.4-0.10 

5-7 

40 

45° 

0.3183 

0.3180 

—  0.0003* 

1-0.25 

0.5-0.10 

4-6 

40 

450 

0.3183 

0.3176 

—  0.0007! 

0.75-0.25 

O  .  4-0  .  TO 

6 

20 

450 

Gauze  cone  used  as  cathode. 


0,2687 

0.2686 

—  O.OOOI* 

i.  5-i 

3-2 

4-6 

25 

500 

0.2687 

o  .  2683 

—0.0004* 

1-25 

2-0.5 

6-7 

30 

450 

0.2687 

o.  2684 

-0.0003* 

i-5 

2-1 

4-6 

25 

450 

0.2687 

0.2686 

—  O.OOOI* 

1-0.25 

2-0.5 

4-6 

25 

450 

0.5375 

0.5373 

—  O.OOO2* 

i  -5-i 

3-2 

6-7 

25 

450 

0.5375 

0.5371 

—  0.0004* 

i  -5-i 

3-2 

6-7 

25 

500 

*  Deposition  complete,  as  shown  by  H2S  test, 
t  Deposition  incomplete,  as  shown  by  HjS  test. 

The  details  of  other  experiments  in  which  the  silver  was  first 
precipitated  as  silver  chloride  and  then  deposited  from  the  solu- 
tion in  ammonia  and  ammonium  oxalate  are  given  in  the  follow- 
ing table.  In  these  experiments,  in  which  the  solutions  were 
more  strongly  ammoniacal  than  those  of  the  .experiments  of  the 
preceding  series,  the  deposits  were  dark  and  spongy,  but  they 
became  lighter  in  color  and  more  compact  upon  drying. 


140  METHODS  IN  CHEMICAL  ANALYSIS 

Electrolysis  of  Silver  Chloride  Dissolved  in  Ammonium  Oxalate  and  Ammonia. 


Current. 

Agin 

AgNO3 
taken. 

Ag  found. 

Error. 

Time. 

tions  per 

Amp. 

Approx. 

N.  D.100. 

Volt. 

mm. 

grm. 

grm. 

grm. 

mm. 

A  crucible  used  as  cathode. 


0.3191 
0.3191 

0.3187 
0.3189 

—0.0004 

—  O.OOO2 

i  -5-i 
i  -5-i 

5-33 
5-33 

5-7 
4-6 

20 
30 

500 
500 

Gauze  disks  used  as  cathode. 


0.3191 

0.3185 

—  0.0006* 

i  -5-i 

0-75-0-5 

5-7 

15 

500 

0.3191 

0.3189 

—  O.OOO2 

i.  5-i 

o.75-o.5 

S-7 

25 

500 

0.3191 

0.3190 

—  o.oooi 

i  -5-i 

0-75-0-5 

4-6 

35 

500 

*  Time  of  electrolysis  short. 

It  appears,  therefore,  that  silver  may  be  deposited  from  an 
ammoniacal  solution  of  the  oxalate,  in  presence  of  ammonium 
nitrate  or  ammonium  chloride,  in  pure  condition  and  in  form 
suitable  for  quantitative  estimation. 

When  the  current  density  is  high  the  deposit  is  apt  to  be 
voluminous,  shrinking  considerably  upon  drying,  and  this  phe- 
nomenon was  especially  notable  in  deposits  upon  the  compara- 
tively small  and  smooth  surface  of  the  crucible.  The  best  form 
of  apparatus  for  this  process  appears  to  be  the  gauze  cone  set 
point  downward,  and  so  placed  with  relation  to  an  annular  plati- 
num band  used  as  the  anode  that  the  end  of  the  axis,  where  the 
mechanical  effect  of  rotation  is  least,  shall  not  receive  much  of 
the  deposit.  Experiments  made  with  stationary  gauze  electrodes 
were  not  successful,  nor  were  those  made  with  a  dish  cathode  and 
stirring  anode. 

The  lodometric  Estimation  of  Silver,  based  upon  the  Use  of 
Potassium  Chromate  as  a  Precipitant. 

With  proper  precautions,  silver  may  be  precipitated  and  accu- 
rately estimated  as  silver  .chromate.*  The  addition  of  a  sufficient 
excess  of  potassium  chromate  to  a  solution  of  silver  nitrate,  even 
iri  the  presence  of  small  amounts  of  nitric  acid,  brings  about 
a  complete  precipitation  of  the  silver  as  silver  chromate,  and 

*  See  page  136. 


COPPER;   SILVER;    GOLD 


141 


the  precipitate  thus  obtained  may  be  transferred  to  the  asbestos 
filter  by  means  of  a  dilute  solution  of  potassium  chromate  and 
washed  with  small  amounts  of  water  without  appreciable  loss  of 
silver  chromate.  Upon  the  basis  of  such  exact  precipitation  of 
silver  chromate  by  potassium  chromate,  Gooch  and  Bosworth* 
have  accomplished  the  iodometric  estimation  of  silver,  both  by 
the  determination  of  the  chromic  acid  ion  of  the  potassium 
chromate  which  remains  after  the  precipitation  of  the  silver  salt 
by  a  known  amount  of  standard  potassium  chromate,  and  by  the 
estimation  of  the  chromic  acid  ion  of  the  precipitated  and  washed 
silver  chromate. 

Precipitation  of  Silver  Chromate  and  Determination  of  the  Excess  of  the  Pre- 
cipitant. 


Silver  taken. 

K2CrO<. 

NaNO3 

present. 

grm. 

Na2S203 
used. 

cm.8. 

Silver 
found. 

grm. 

Error  in 
terms  of 
silver. 

grm. 

Volume  of 
solution. 

cm.3 

Weight, 
grm. 

Volume. 
cm.s. 

Weight, 
grm. 

20 

O.I26l 

0.3039 

26.54 

o.  1262 

+O.OOOI 

25 

0.1576 

0.3039 

22.61 

0.1575 

—  o.oooi 

15 

o  .  0946 

0.3293 

30-5I 

o  .  0946 

o.oooo 

IS 

o  .  0946 

0.3293 

30.60 

o  .  0940 

—0.0006 

15 

o  .  0946 

0.3293 

30.57 

0.0941 

—0.0005 

15 

o  .  0946 

0.3293 

30.55 

0.0945 

—  O.OOOI 

19.98 

o.  1260 

0.3293 

32.20 

0.1255 

-0.0005 

2O 

0.1261 

0.3293 

32-09 

o  .  i  263 

+O.OOO2 

2O 

o.  1261 

0.3293 

32.10 

0.1263 

+O.OOO2 

25 

0.1576 

0.3293 

27.30 

0.1576 

o.oooo 

10 

O.  IIOI 

37 

0.2436 

20.47 

o.  1107 

+0.0006 

IO 

O.IIOI 

37 

0.2436 

20.53 

0.1103 

+O.OOO2 

IO 

O.IIOI 

25 

0.1647 

9.07 

0.1097 

—0.0004 

IO 

O.IIOI 

27 

0.1778 

1  1.  02 

o.  1096 

-0.0005 

15 

0.0862 

30 

0.1974 

17.14 

0.0859 

—0.0003 

15 

0.0862 

3° 

0.1974 

17-05 

0.0865 

+0.0003 

15 

0.0862 

30 

0.1974 

17.14 

0.0859 

—0.0003 

25 

0.1437 

50 

0.3294 

10 

28.53 

0.1435 

—  O.OOO2 

According  to  the  first  procedure  a  known  amount  of  standard 
potassium  chromate  in  excess  of  the  amount  needed  to  precipi- 
tate the  silver  is  added  to  the  solution  of  silver  nitrate.  The  pre- 
cipitate is  dissolved  in  ammonia  and  reprecipitation  brought  about 
by  boiling  to  a  volume  of  10  cm.3-i5  cm.3.  The  second,  crystal- 
line precipitate  is  filtered  upon  asbestos  and  washed  with  the  least 
possible  amount  of  water  applied  in  small  portions  successively. 
The  filtrate  is  treated  with  potassium  iodide  and  acidified  with 

*  F.  A.  Gooch  and  Rowland  S.  Bosworth,  Am.  Jour.  Sci.,  [4],  xxvii,  302. 


142 


METHODS  IN  CHEMICAL  ANALYSIS 


sulphuric  acid.  The  iodine  set  free  is  titrated  with  sodium  thio- 
sulphate.  The  difference  between  the  silver  value  of  the  iodine 
thus  found  and  that  of  the  potassium  dichromate  used  is  taken 
as  the  measure  of  the  silver  present.  In  the  table  are  given 
the  details  of  experiments  performed  in  accordance  with  this 
procedure. 

Inasmuch  as  relatively  large  amounts  of  potassium  chromate 
are  necessary  to  bring  about  complete  precipitation  of  the  silver 
chromate  in  the  presence  of  nitric  acid,  the  above  procedure 
is  less  adapted  to  the  determination  of  silver  in  a  solution  con- 
taining that  acid  than  the  second  method  whereby  the  precipi- 
tated and  washed  silver  chromate  is  determined.  In  this  second 
method  the  silver  solution  containing  free  nitric  acid  is  treated 
with  potassium  chromate  in  excess  of  the  amount  necessary  to 
take  up  the  nitric  acid  with  formation  of  potassium  dichromate. 
The  precipitate  is  dissolved  in  ammonia  and  reprecipitation  is 
effected  by  boiling  to  a  volume  of  10  cm.3-i5  cm.3.  The  second, 
crystalline  precipitate  is  transferred  to  an  asbestos  filter  by 
means  of  a  dilute  solution  of  potassium  chromate,  washed  with 
the  least  possible  amount  of  water  applied  in  small  portions 
successively,  and  dissolved  in  a  few  cm.3  of  a  strong  solution  of 
potassium  iodide.  The  solution  in  potassium  iodide  is  diluted 
and  acidified  with  sulphuric  acid.  The  iodine  set  free  is  titrated 
with  sodium  thiosulphate  and  taken  as  the  measure  of  the  silver 
present. 

Results  of  this  procedure  are  given  in  the  following  table. 

lodometric  Determination  of  Precipitated  Silver  Chromate. 


Silver  taken. 

HNO3 
present. 

K2CrO4, 
weight. 

Na2S203 
used. 

Silver 
found. 

Error  in 
terms  of 
silver. 

Volume  of 
solution. 

Weight. 

cm.* 

grrn. 

grm. 

grm. 

cm.3 

grm. 

grm. 

25 

0.1348 

0.063 

O.6o 

lS-35 

0.1342 

—  O.O006 

2O 

o.  1078 

0.063 

0.65 

14.67 

0.1073 

—  0  .  0005 

IS 

o  .  0808 

0.063 

0.65 

II  .06 

o  .  0809 

+  O.OOOI 

15 

0.0808 

0.063 

0.65 

10.96 

o  .  0802 

—  O.OOO6 

20 

o.  1078 

0.063 

0.65 

14.67 

0.1073 

-0.0005 

3° 

0.1618 

0.063 

0-75 

22.02 

0.1610 

—  0.0008 

20 

0.1078 

0.063 

0.65 

14.71 

0.1075 

—0.0003 

25 

0.1348 

0.063 

0.65 

I8.4I 

o.i347 

—  o.oooi 

25 

0-1348 

0.063 

0.65 

I8.4I 

o.i347 

—  O.OOOI 

20 

0.1078 

0.063 

0.65 

14-74 

o.  1078 

O  .  0000 

COPPER;   SILVER;   GOLD 


143 


The  lodometric  Determination  of  Silver  Based  upon  the  Reducing 
Action  of  Potassium  Ar senile. 

It  is  well  known  that  an  ammoniacal  solution  of  silver  arsenite 
deposits  metallic  silver  when  the  ammonia  is  evaporated  by  boil- 
ing. During  this  reaction,  by  which  the  silver  salt  is  reduced, 
the  arsenious  acid  becomes  oxidized  to  the  higher  condition  of 
oxidation,  where  arsenic  has  a  valence  of  five,  according  to  the 
equation 

2  Ag2O  +  As2O3  =  As2O5  +  4  Ag. 

Bosworth*  has  shown  that  this  reaction, is  quantitative  and 
capable  of  serving  as  the  basis  of  a  reliable  iodometric  method 
for  the  determination  of  silver.  To  the  solution  of  silver  taken 
as  the  nitrate  is  added  a  known  volume  of  a  standard  potassium 
arsenite  solution  in  excess  of  the  amount  necessary  to  reduce  the 
silver  salt  present.  Ammonia  is  added  in  sufficient  quantity  to 
dissolve  the  precipitate  formed,  or  25  cm.3  of  a  saturated  solu- 
tion of  sodium  bicarbonate  may  be  used  instead  of  the  ammo- 
nia. The  resulting  solution  is  diluted  to  100  cm.3,  and  boiled. 
The  solution,  out  of  which  metallic  silver  separates,  is  cooled, 
slightly  acidified,  and  then  made  alkaline  with  sodium  bicarbo- 
nate. The  excess  of  potassium  arsenite  is  titrated  with  n/io 
iodine.  The  silver  value  of  the  iodine  used  is  subtracted  from 
that  of  the  potassium  arsenite  originally  taken,  and  the  result 
used  as  a  measure  of  the  silver  present. 

Results  of  this  procedure  are  given  in  the  table. 


Reduction  of  Silver  by  Arsenite  and  Titration  of  the  Excess. 


Silver  taken, 
grrn. 

KH2AsO3  added. 

I  used. 

Silver  found, 
grm. 

Error  in  terms 
of  silver. 

gnu. 

cm.3 

Silver  value, 
grrn. 

cm.3 

Silver  value, 
grm. 

Use  of  NH4OH  with  filtering. 


0.1054 

20 

O.  2OOO 

8-53 

0.0949 

0.1051 

—0.0003 

0.1054 

20 

O  .  2OOO 

8.52 

o  .  0948 

0.1052 

—  O.OOO2 

0.1054 

30 

0.3000 

17.42 

0.1939 

o.  1061 

+o  .  0007 

0.1159 

2O 

0.2000 

7.60 

o  .  0846 

0.1154 

—  0.0005 

0.1054 

21 

O.2IOO 

9-37 

0.1043 

0.1057 

+0.0003 

0.1054 

2O 

O.  2OOO 

8.48 

0.0944 

o.  1056 

+0  .  0002 

*  Rowland  S.  Bosworth,  Am.  Jour.  Sci.  [4],  xxviii,  287. 


144 


METHODS  IN  CHEMICAL  ANALYSIS 


Reduction  of  Silver  by  Arsenite  and  Titration  of  the  Excess. 


Silver  taken, 
grm. 

KH2AS03  added. 

I  used. 

Silver  found, 
grm. 

3rror  in  terms 
of  silver. 

grm. 

cm.3 

Silver  value, 
grm. 

cm.3 

Silver  value, 
grm. 

Use  of  NaHCO3  with  filtering. 


0.1054 

15 

0.1618 

5.65 

0.0571 

0.1047 

—  0.0007 

0.1054 

23 

0.2481 

14.14 

0.1430 

0.1051 

-0.0003 

o.  1054 

12 

0.1295 

2.40 

0.0243 

0.1052 

—  O.OO02 

0.1054 

15 

0.1618 

5.60 

0.0566 

0.1052 

—  O.O002 

0.1054 

15 

0.1618 

5-55 

0.0561 

0.1057 

+0.0003 

0.1054 

20 

0.2158 

10.91 

0.1104 

0.1054 

o.oooo 

0.2635 

35 

0.3776 

n-33 

0.1146 

0.2630 

—0.0005 

Use  of  NH4OH.     Titration  carried  on  in  presence  of  the  precipitate. 

0.1054 

20 

O.2OOO 

8-55 

0.0952 

0.1048 

—  0.0006 

0.1054 

20 

O.2OOO 

8.50 

0.0946 

0.1054 

o.oooo 

0.1054 

23 

0.2300 

11.28 

0.1256 

o.  1044 

—  O.OOIO 

0.1054 

2O 

O.2OOO 

8-45 

0.0941 

0.1059 

+0.0005 

0.1054 

2O 

O.2OOO 

8.48 

0.0944 

o.  1056 

+O.OOO2 

Use  of  NaHCO3.     Titration  carried  on  in  presence  of  the  precipitate. 


0.1054 

18 

o.  1800 

6.80 

0.0757 

0.1043 

—  o.oon 

0.1054 

17 

O.I7OO 

5-8l 

0.0647 

0.1053 

—  o.oooi 

0.1054 

15 

0.1500 

4.00 

o  .  0445 

0-1055 

+O.OOOI 

0.1054 

21 

O.2IOO 

9-45 

0.1052 

0.1048 

—0.0006 

0.1054 

25 

0.2500 

13.00 

0.1447 

0.1053 

—  O.OOOI 

0.1054 

31 

0.3100 

18.40 

o  .  2048 

0.1052 

—  O.OOO2 

Use  of  NaHCO3.     2  grm.  of  NaNO3  present.     Titration  carried  on  in 
presence  of  precipitate. 


0.0949 

21 

0.2100 

IO.42 

0.1160 

o  .  0940 

—0.0009 

0.1054 

21 

O.  2IOO 

9-43 

o.  1050 

o.  1050 

—0.0004 

0.1265 

2O 

O.  2OOO 

6.60 

0.0735 

0.1265 

o.oooo 

0.1686 

21 

O.2IOO 

3.8o 

0.0432 

0.1678 

—0.0008 

0.1054 

15 

0.1500 

4.08 

0.0454 

0.1046 

—0.0008 

This  process  proves  to  be  applicable  to  the  determination  of 
the  silver  in  freshly  precipitated  silver  chloride,  so  that  its  range 
is  thus  extended  to  the  determination  of  silver  in  many  mixtures. 

According  to  the  procedure  outlined,  the  freshly  precipitated 
silver  chloride  is  acted  upon  by  ammonia  until  dissolved.  The 


COPPER;   SILVER;    GOLD 


solution  is  diluted  to  100  cm.3  and  the  reduction  is  effected, 
adding  an  excess  of  standard  arsenite  and  boiling  the  mixture. 
The  excess  of  arsenite  is  titrated  according  to  the  method  de- 
scribed above. 

Results  of  this  procedure,  including  separations  from  copper 
and  lead,  are  given  in  the  following  table. 

Reduction  of  Silver  Chloride  after  Separations. 


Silver  taken. 

grm. 

KH2AsO3  added. 

I  used. 

Silver  found, 
gnu. 

Error  in  terms 
of  silver. 

grui. 

cm.» 

Silver  value. 

grm. 

cm.8 

Silver  value, 
grm. 

Reduction  of  precipitated  AgCl. 


0.1017 

is 

0.1619 

5-4° 

0.0599 

O.IO2O 

+0.0003 

0.1017 

15 

o.  1619 

5-44 

0.0603 

o.  1016 

—  O.OOOI 

0.1017 

15 

0.1619 

5-40 

0.0599 

0.1020 

+0.0003 

0.1017 

15 

o.  1619 

5-42 

o  .  0601 

0.1018 

+0.0001 

0.1017 

17 

0.1834  • 

7-44 

0.0825 

0.1009 

—0.0008 

Reduction  of  AgCl  precipitated  in  the  presence  of  0.09  grm.  of  copper. 


0.1017 

15 

0.1619 

5-41 

0.0600 

o.  1019 

+O.OOO2 

0.1017 

is 

o.  1619 

5-44 

0.0603 

0.1016 

—  O.OOOI 

0.1017 

is 

0.1619 

5-39 

0.0598 

O.IO2I 

+0.0004 

Reduction  of  AgCl  precipitated  in  the  presence  of  0.2  grm.  of  lead. 


O.  I22O 

o.  1108 

16 

15 

0.1726 
0.1619 

4.57 

4.60 

0.0507 
0.0510 

0.1219 
0.1109 

—  O.OOOI 
+  0.0001 

Reduction  of  AgCl  precipitated  from  a  solution  containing  0.09  grm.  of 
copper  and  0.2  grm.  of  lead. 


0.1017 

15 

0.1619 

5-45 

o  .  0604 

0.1015 

—  O.OOO2 

GOLD. 
The  Electrolytic  Determination  of  Gold. 

Medway*  has  shown  that  from  an  ammoniacal  cyanide  solu- 
tion gold  may  be  deposited  rapidly  and  in  good  form  upon  the 
rotating  crucible  used  as  the  cathode. 

*  H.  E.  Medway,  Am.  Jour.  Sci.,  [4],  xviii,  56. 


146 


METHODS  IN  CHEMICAL  ANALYSIS 


Deposition  from  an  Ammoniacal  Cyanide  Solution. 


Gold  taken. 

Gold  found. 

Error. 

Current. 

N.  D.100. 

Time. 

grm. 

grm. 

grm. 

amp. 

mm. 

0.0695 

o  .  0694 

—  0.0001 

2 

6.6 

3° 

0.0695 

o  .  0696 

+O.OOOI 

2           * 

6.6 

3° 

0.0598 

0.0598 

o.oooo 

0-5 

1.8 

3° 

0.0598 

0.0598 

o.oooo 

0-5 

1.8 

3° 

0.0598 

0-05975 

—0.00005 

I 

3-3 

25 

The  lodometric  Estimation  of  Small  Amounts  of  Gold. 
Gooch  and  Morley*  have  shown  that  when  potassium  iodide 
reacts  at  suitable  concentration  upon  small  amounts  of  gold  tri- 
chloride in  solution,  the  reaction  takes  place  regularly  and  in 
accordance  with  the  theory  that  two  molecules  of  the  thiosul- 
phate  are  the  equivalent  of  two  atoms  of  iodine  and  one  atom 
of  gold. 
AuCl3  +  2  KI  +  2  Na2S2O3  =  AuCl  +  Na2S4O6  +  2  KC1  +  2  Nal. 

The  reduction  of  the  auric  salt,  with  the  consequent  liberation 
of  iodine,  is,  however,  conditioned  by  the  volume  of  the  solution, 
the  mass  of  the  iodine  present,  and  the  time  of  action.  The 
following  statement,  in  which  each  result  is  the  average  of 
several  titrations  in  close  agreement,  shows  the  effect  upon  the 
immediate  evolution  of  iodine  brought  about  by  adding  varying 
amounts  of  water  to  the  gold  solution  before  introducing  the 
iodide,  and  the  effect  of  different  amounts  of  iodide  at  different 
dilutions. 


Volume 

Potassium  iodide. 

Gold 
chloride. 

before  the 
addition  of 
the  thio- 

sulphate. 

o.oi  grm. 

0.02  grm. 

0.05  grm. 

o.i  grm. 

0.2  grm. 

grm. 

cm. 

£8 

0.8l 

0.81 

0.81 

0.82 

0.84 

0.00087 

15 

S  Jc\ 

0.77 

0.78 

0.80 

0.81 

0.8l 

11 

25 

.3  ^  8  g 

0.74 

0.72 

0.78 

0.79 

0.80 

50 

'Q    w  ~    § 

0.61 

0.61 

0.68 

o.  76 

0.79 

" 

IOO 

w  is  2 

0-45 

0.49 

0.60 

0.72 

0-75 

" 

2OO 

"5  c 

It  is  evident  that  for  the  smaller  amounts  of  iodide  the  libera- 
tion of  iodine  decreases  rapidly  with  the  dilution.     The  larger 
amounts  at  the  highest  concentration  show  readings  a  trifle 
*  F.  A.  Gooch  and  Frederick  H.  Morley,  Am.  Jour.  Sci.,  [4],  viii,  261. 


COPPER;   SILVER;   GOLD 


147 


above  the  normal  —  perhaps  because  the  well-known  effect  of 
concentrated  solutions  of  a  soluble  iodide  upon  the  delicacy  of 
the  starch  end-color  begins  to  appear.  At  volumes  lying  be- 
tween the  limit  of  25  cm.3  and  50  cm.3,  o.i  grm.  of  potassium 
iodide  is  an  appropriate  amount  to  use;  at  a  volume  of  15  cm.3, 
o.oi  grm.  to  0.05  grm.  of  the  iodide  will  do  the  work;  and  at 
lower  dilutions,  as  will  appear  in  the  tabular  statements  to  follow, 
even  less  of  the  iodide  is  effective. 

In  carrying  out  the  process  of  analysis,  a  convenient  amount 
of  the  solution  of  gold  chloride  is  drawn  from  a  burette,  potassium 
iodide  is  introduced  in  amounts  always  several  times  the  theoret- 
ical equivalent  of  the  gold,  and  more  than  enough  to  dissolve  the 
aurous  iodide  precipitated  at  first,  a  sufficiency  of  clear  starch 
indicator  is  added,  the  starch  blue  bleached  by  standardized 
thiosulphate,  and  standardized  iodine  added  until  the  liquid 
assumes  a  faint,  rose  color. 

Experimental  results  are  given  in  the  following  table: 

Solutions  Approximately  n/ioo 

Gold  chloride  =0.8710  grm.  to  i  liter. 

Sodium  thiosulphate,  nearly  n/ioo,  =1.7012  grm.  to  i  liter. 
Iodine,  nearly  w/ioo,  =1.3697  grm.  to  i  liter. 

Volume  at  beginning  of  titration,  approximately  50  cm.3. 


AuCl3 
taken. 

cm.3 

KI  taken. 
grm. 

Na2S2O3  used. 
cm.3 

Gold  found, 
grm. 

Theory  for 
gold. 

grm. 

Error, 
grm. 

5 

0.05 

4.02 

0.00426 

0.00435 

—  0.00009 

5 

0.05 

4.01 

0.00425 

0.00435 

—  o.oooio 

5 

0.05 

4.06 

0.00431 

0.00435 

—0.00004 

5 

0.05 

4.07 

0.00432 

0.00435 

—0.00003 

5 

0.05 

4-*04 

0.00428 

0.00435 

—  0.00007 

10 

0.08 

8.17 

0.00867 

0.00871 

—  0.00004 

10 

0.08 

8.15 

o  .  00864 

0.00871 

—  0.00007 

10 

O.o8 

8.16 

o  .  00865 

0.00871 

—0.00006 

10 

O.o8 

8.15 

o  .  00864 

0.00871 

—  0.00007 

10 

O.o8 

8.19 

o  .  00869 

0.00871 

—  O.OOOO2 

10 

O.o8 

8.46 

0.00897 

0.00871 

+0  .  00026 

10 

0.08 

8.24 

o  .  00874 

0.00871 

-j-o  .  00003 

When  approximately  centinormal  solutions  of  gold,  iodine  and 
thiosulphate  are  used,  an  error  of  o.oi  cm.3  in  reading  the  volume 
corresponds  to  an  error  of  o.ooooi  grm.  of  gold.  It  is  not  to  be 
expected  that  such  readings  can  be  trusted  ordinarily  to  a  higher 
degree  of  accuracy  than  0.02  cm.3.  In  case  all  three  solutions 
should  be  read  to  this  limit  of  accuracy  with  the  errors  of  all 


148 


METHODS  IN  CHEMICAL  ANALYSIS 


lying  in  the  same  direction,  the  summation  of  error  would  corre- 
spond to  0.00006  grm.  of  gold. 

Errors  in  reading  are,  of  course,  reduced  when  n/iooo  solu- 
tions are  employed,  but  the  use  of  n/iooo  iodine  necessitates 
a  correction  of  o.i  cm.3  for  volumes  not  exceeding  30  cm.3,  that 
being  the  amount  necessary  to  bring  out  the  rose  color  in  blank 
tests  containing  no  gold.  After  the  introduction  of  o.i  cm.3  of 
n/iooo  iodine  into  a  mixture  of  potassium  iodide  and  starch 
indicator  of  volume  not  exceeding  30  cm.3,  a  single  drop  of  the 
gold  solution  —  equivalent  to  0.000002  grm.  of  gold  —  gives  a 
distinct  rose  color;  before  such  adjustment  of  the  solution  five 
drops  —  equivalent  to  o.ooooio  grm.  of  gold  —  must  be  added 
to  develop  the  same  color. 

The  following  table  gives  the  data  of  tests  with  such  solutions. 

Solutions  Approximately  n/iooo 

Gold  chloride  =0.0871    grm.  to  i  liter. 

Sodium  thiosulphate,  nearly  w/iooo    =0.17012  grm.  to  i  liter. 
Iodine,  nearly  w/iooo  =0.13697  grm.  to  i  liter. 


AuClj 
taken. 

cm.s 

KI  taken, 
grm. 

Na2S2O3  used. 
cm.3 

Gold  taken, 
grm. 

Gold  found, 
grm. 

Error, 
grm. 

10 

O.OI 

8-39 

0.000871 

o  .  000890 

+0.000019 

.9 

O.OI 

7-45 

0.000784 

0.000790 

+0  .  000006 

8 

O.OI 

6.30 

0.000697 

0.000668 

—  0.000029 

7 

0.008 

5-So 

O.OOo6lO 

0.000583 

—  0.000027 

6 

0.008 

5-12 

0.000523 

o  .  000543 

+  O.OOOO20 

5 

0.005 

4-23 

0.000435 

o  .  000449 

+0.000014 

4 

0.005 

3.38 

o  .  000348 

0.000358 

+O.OOOOIO 

3 

0.003 

2-55 

0.000261 

O.OOO27O 

+0.000009 

2 

0.003 

1.71 

0.000174 

0.000181 

+0.000007 

I 

0.003 

0.90 

0.000087 

0.000095 

+0.000008 

In  the  practical  application  of  the  process  to  the  determination 
of  gold,  the  elementary  form  of  that  metal  is  the  natural  starting 
point.  To  get  the  metal  into  solution  with  chlorine  water  or 
mixed  hydrochloric  and  nitric  acids  is  an  easy  matter,  but  the 
removal  of  the  excess  of  the  oxidizer  by  evaporation  without 
reducing  some  auric  chloride  to  the  aurous  form  is  difficult. 
Free  chlorine  may,  however,  be  removed  from  a  solution  of  auric 
chloride,  without  reducing  the  auric  salt,  by  treatment  of  the 
solution  with  ammonia  in  excess,  boiling  gently,  acidifying  with 
hydrochloric  acid,  and  heating  if  necessary  to  redissolve  the  pre- 
cipitate by  ammonia,  again  treating  with  ammonia  and  heating. 


COPPER;   SILVER;   GOLD 


149 


and  once  more  acidifying.  On  the  second  addition  of  ammonia 
no  precipitation  usually  takes  place  with  these  small  amounts 
of  gold. 

The  following  table  contains  determinations  made  with  a  solu- 
tion of  pure  gold  leaf. 

lodometric  Determination  of  Gold, 

Gold  chloride  made  by  dissolving  0.0104  grm-  of  pure  gold  in  the  manner 
described  and  diluting  to  200  cm.3. 

Sodium  thiosulphate,  nearly  w/iooo,    =0.17012  grm.  to  i  liter. 
Iodine,  nearly  w/iooo  =0.13697  grm.  to  i  liter. 

Potassium  iodide  =iogrm.  to  i  liter. 

Portions  were  treated  with  the  potassium  iodide  without  previous  dilu- 
tion. 


AuClj 
taken. 

cm.» 

KI  taken, 
grm. 

Na2S2O3  used. 
cm.3 

Gold  taken, 
grm. 

Gold  found, 
grm. 

Error, 
grm. 

I 

0.005 

o-55 

0.000052 

0.000058 

+0  .  000006 

I 

0.005 

0-55 

0.000052 

0.000058 

+0.000006 

2 

0.005 

i.  06 

O.OOOIO4 

O.OOOII2 

-J-O.OOOOOS 

2 

0.005 

1.08 

O.OOOIO4 

O.OOOII4 

+O.OOOOIO 

5 

O.OI 

2.45 

0.000260 

O.OOO26O 

O  .  OOOOOO 

5 

O.OI 

2.50 

O.OOO26o 

O.OOO265 

+o  .  000005 

5 

O.OI 

2-45 

o  .  000260 

O.OOO26O 

O  .  OOOOOO 

5 

O.OI 

2.50 

O.OOO26O 

O.OOO265 

+0.000005 

5 

O.OI 

2.50 

0.000260 

0.000265 

+0.000005 

10 

O.O2 

4.86 

0.000520 

0.000515 

—  0.000005 

10 

O.O2 

4-85 

0.000520 

0.000517 

—0.000003 

IO 

O.O2 

4.90 

0.000520 

O.OOO52O 

0  .  OOOOOO 

IO 

O.O2 

4.80 

0.000520 

O.OOO5I2 

—0.000008 

IO 

O.O2 

4.84 

0.000520 

O.OOO5I6 

—0.000004 

To  show  the  range  and  error  of  the  process,  the  results  of  these 
and  other  experiments  recorded  by  Gooch  and  Morley  may  be 
summarized  as  follows: 

Range  of  the  Process. 


Strength  of  solutions. 

Number  of 
determina- 
tions. 

Gold  taken. 

Error, 
average. 

Extremes. 

Iodine. 

Thiosul- 
phate. 

Gold  in 
in  i  cm.3 

rng. 

mg. 

mg. 

mg. 

II 

8.71   -4-35 

n/ioo 

W/IOO 

0.871 

—0.05 

(  +0.03 
(  —O.I 

20 

0.871-0.087 

n/ioo 

n/ioo 

0.0871 

+0.02 

(  +0.06 
(    —  O.O2 

IO 

0.871-0.087 

n/iooo 

W/IOOO 

0.0871 

+0.004 

(   +O.O2O 
(   —O.O29 

14 

0.520-0.052 

n/iooo 

tt/IOOO 

0.052 

+0.002 

(  +0.01 
{    —0.008 

150  METHODS  IN  CHEMICAL  ANALYSIS 

It  is  plain  that  the  average  experimental  errors,  due  to  all 
causes,  do  not  very  much  exceed  the  errors  which  might  natu- 
rally be  expected  to  arise  from  errors  of  reading. 

In  repeating  this  work,  Maxson  *  has  obtained  results  of  a 
similar  order  of  accuracy.  Maxson  has  also  studied  the  possi- 
bility of  reduction  of  the  aurous  iodide  formed  in  the  process  and 
finds  in  periods  much  exceeding  those  required  for  the  analytical 
operation  no  evidence  of  further  action  other  than  the  for- 
mation of  the  aurous  salt.  Thus,  aurous  iodide,  obtained  by 
treating  a  solution  of  auric  chloride,  containing  0.0125  grm.  of 
gold,  with  potassium  iodide  according  to  the  directions  of  Gooch 
and  Morley,  adding  starch  and  bleaching  the  starch  iodide  with 
sodium  thiosulphate,  shows  no  color  of  starch  blue  after  the  in- 
terval of  an  hour.  Inasmuch  as  an  interval  of  ten  minutes  is 
enough  for  the  complete  manipulation  of  a  single  determination, 
it  is  plain  that  the  stability  of  the  aurous  iodide  does  not  figure 
in  the  accuracy  of  the  determination  of  the  small  amounts  of 
gold  for  which  the  process  was  designed. 

The  Colorimetric  Determination  of  Small  Amounts  of  Gold. 

A  colorimetric  method  for  the  estimation  of  small  amounts  of 
gold  has  been  based  by  Maxson  f  upon  the  coloration  exhibited 
by  suspensions  of  red  colloidal  gold. 

Blake  J  has  shown  that  acetylene  is  the  most  suitable  reagent 
for  effecting  the  reduction  of  the  auric  salt  with  production  of 
the  red  colloidal  gold,  the  treatment  consisting  in  drying  the 
chloride  at  170°,  dissolving  in  ether,  and  pouring  the  etheral 
solution  into  water  containing  ether  and  saturated  with 
acetylene  gas.  For  the  purposes  of  this  method  the  simpler 
treatment  with  an  aqueous  solution  of  acetylene  proved  to  be 
adequate.  The  procedure  consists  in  preparing  a  red  colloidal 
suspension  containing  a  known  amount  of  gold  in  a  given  volume 
and  by  means  of  measured  amounts  of  this  solution  matching 
the  color  developed  in  a  similar  solution  containing  in  measured 
volume  the  gold  to  be  determined. 

For  making  the  comparisons  of  color  a  modified  form  of  the 
apparatus  proposed  by  Penfield  for  the  colorimetric  estimation 

*  Ralph  N.  Maxson,  Am.  Jour.  Sci.,  [4],  xvi,  155;  xvii,  466. 
t  Ralph  Nelson  Maxson,  Am.  Jour.  Sci.,  [4]  xxi,  270. 
J  Am.  Jour.  Sci.,  [4],  xvi,  381. 


COPPER;  SILVER;  GOLD  151 

of  titanium  is  used.  This  consists  of  comparison  tubes  set  ver- 
tically in  a  dark  box  and  illuminated  from  below.  A  mirror  suit- 
ably situated  beneath  the  box  containing  the  tubes  gives  efficient 
illumination.  Such  an  apparatus  is  cheaply  and  easily  procured. 
With  tubes  having  a  diameter  of  I  cm.  and  a  length  of  13  cm., 
and  accurately  graduated  to  hold  10  cm.3,  this  simple  colorimeter 
is  capable  of  determining  accurately  very  small  amounts  of  gold. 

The  standard  suspension  is  made  by  treating  in  a  measuring 
flask  a  measured  amount  of  a  standardized  solution  of  ordinary 
undried  gold  chloride  with  an  aqueous  solution  of  acetylene  and, 
after  the  full  development  of  color,  making  up  to  the  mark. 

A  solution  of  pure  gold  to  be  examined  is  treated  similarly  with 
aqueous  acetylene  and  made  up  to  known  volume.  When  small 
amounts  are  handled  the  volume  of  the  solution  should  not  exceed 
a  few  cubic  centimeters,  and  only  a  small  amount  of  the  aqueous 
solution  of  acetylene  should  be  added;  otherwise  the  coloration 
may  be  partially  or  totally  inhibited. 

If  traces  of  electrolyte  are  present,  the  coagulation  of  the  red 
gold  may  sometimes  be  avoided  by  the  addition  of  a  few  drops 
of  ether  to  the  cold  solution. 

It  is  a  well-known  fact  that  small  amounts  of  electrolyte  will 
rapidly  change  red  gold  to  the  blue  modification.  It  is  necessary, 
therefore,  to  conduct  the  comparisons  in  a  room  reasonably  free 
from  fumes,  and  to  have  all  containing  vessels  free  from  soluble 
material.  Flasks  which  have  been  treated  with  steam  for  a  few 
minutes  give  the  best  results.  Red  suspensions  contained  in 
such  flasks  may  show  no  trace  of  blue  after  an  interval  of  several 
weeks. 

The  comparison  of  colors  is  carried  out  in  the  following  man- 
ner: A  measured  amount  of  the  suspension  is  drawn  off  into 
the  left-hand  tube  and  diluted  to  the  mark  with  water ;  a  suitable 
amount  of  water  is  then  placed  in  the  right-hand  tube  and  the 
standard  suspension  drawn  into  the  tube  until  the  colors  are 
seen  to  be  identical.  The  amount  of  water  to  be  used  can 
be  determined  by  preliminary  experiment.  The  positions  of  the 
tubes  are  reversed  before  the  final  reading,  and  the  mean  taken. 

In  the  table  below  is  the  record  of  experiments  made  to 
determine  the  range  of  amounts  of  gold  capable  of  accurate 
estimation  in  such  an  apparatus  with  tubes  of  the  dimensions 
described  above.  The  comparisons  were  made  with  a  red  sus- 


152 


METHODS  IN  CHEMICAL  ANALYSIS 


pension,  prepared  by  careful  dilution  of  a  more  concentrated 
standard  suspension,  which  contained  0.0000107  grm.  of  metal 
in  i  cm.3. 

Colorimetric  Estimation  of  Gold. 


Gold  suspension 
taken. 

cm.* 

Gold  suspension 
used. 

cm.3 

Gold  used, 
gnu  . 

Gold  found. 

grm. 

Error. 

grm. 

9-50 
8.00 
7.00 
6.00 
5.00 

9-05 
7-59 
6.89 

5-83 
4.84 

0.000102 
O.OOOO86 
0.000075 
O.OOOO65 

o  .  000054 

0.000097 
0.000082 
0.000074 

o  .  000063 
0.000052 

—  0.000005 
—0.000004 

—  O.OOOOOI 
—  O.OOOO02 
—  O.OOOOO2 

4.00 

3-88 

o  .  000043 

o  .  000042 

—  O.OOOOOI 

3.00 

2.OO 

2-47 
1.82 

0.000032 

O.OOOO22 

0.000027 

O.OOO02O 

—  0.000005 

—  O.OOOOO2 

I  .OO 

0-93 

O.OOOOII 

O.OOOOIO 

—  O.OOOOOI 

The  intensity  of  the  color  in  the  experiments  ranged  from  a 
deep  red  to  a  faint  pink.  Further  comparisons  made  with  sus- 
pensions more  dilute  than  those  described  above  gave  errors  of 
magnitude  increasing  with  the  dilution.  The  amounts  handled 
here  are,  then,  the  minimum  quantities  that  can  be  accurately 
estimated  with  the  apparatus  described.  It  is  obvious  that  if 
larger  amounts  of  the  metal  are  to  be  determined,  tubes  of  greater 
dimensions  should  be  used. 

The  application  of  such  a  method  for  the  determination  of 
gold  naturally  starts  with  that  element.  The  weighed  amount 
of  metal,  contained  in  a  clean  porcelain  crucible,  can  be  readily 
brought  into  solution  with  the  aid  of  chlorine  water  or  aqua 
regia  and  the  excess  of  the  solvent  evaporated  off  upon  the  water 
bath. 


CHAPTER  IV. 

BERYLLIUM;  MAGNESIUM;  CALCIUM;  STRONTIUM;  BARIUM. 

BERYLLIUM. 

Ammonium  Beryllium  Phosphate. 

BERYLLIUM,  like  calcium,  strontium  and  barium,  cannot  be  es- 
timated by  precipitation  as  a  double  ammonium  phosphate  and 
subsequent  ignition,  as  are  magnesium,  zinc  and  cadmium.  The 
precipitated  ammonium  beryllium  phosphate,  as  has  been  shown 
by  Austin,*  always  contains  triberyllium  phosphate. 

The  Conversion  of  Beryllium  Chloride  to  Beryllium  Oxide. 

Havens  f  has  shown  that  small  amounts  of  beryllium  chloride 
may  be  easily  converted  to  the  oxide,  without  precipitation  and 
filtration,  by  treatment  with  nitric  acid  and  ignition. 

The  solution  of  the  chloride  is  evaporated  just  to  dryness  on 
a  radiator,  care  being  taken  not  to  heat  to  the  volatilizing  point 
of  the  salt,  a  few  drops  of  concentrated  nitric  acid  are  added, 
the  liquid  is  evaporated,  and  the  residue  heated  gently  at  first 
and  finally  in  the  flame  of  the  blast  lamp.  This  conversion  of 
beryllium  chloride  to  beryllium  nitrate  may  be  carried  on  in 
platinum  without  attacking  that  metal  appreciably,  provided 
care  be  taken  to  remove  all  excess  of  hydrochloric  acid  and  to 
add  the  nitric  acid  to  the  dry  residue. 

Conversion  of  Beryllium  Chloride  to  the  Oxide. 


BeO  taken  in 
solution  as  BeCla. 

BeO  found. 

Error. 

grin. 

grm. 

gnu* 

0.0483 
o  .  0483 

0.1076 

0.0481 
0.0483 
0.1085 

—0.0002 

o.oooo 
+0.0009 

*  Martha  Austin,  Am.  Jour.  Sci.,  [4],  viii,  207. 
t  F.  S.  Havens,  Am.  Jour.  Sci.,  [4],  iv,  112. 
153 


154  METHODS  IN  CHEMICAL  ANALYSIS 

The  Separation  of  Beryllium  Oxide  from  Ferric  Oxide. 

Havens  and  Way*  have  separated  beryllium  oxide  from  ferric 
oxide  by  volatilization  of  the  latter  in  a  stream  of  gaseous  hydro- 
gen chloride  charged  with  a  little  free  chlorine,  with  care  to  avoid 
mechanical  loss  through  too  rapid  volatilization  of  the  iron.f 

MAGNESIUM. 

The  Determination  of  Magnesium  by  Precipitation  and  Ignition 
of  Ammonium  Magnesium  Carbonate. 

According  to  Schaffgotsch,J  the  very  concentrated  solution  of 
the  sulphates,  nitrates  or  chlorides  of  magnesium,  sodium  and 
potassium  is  treated  with  a  concentrated  solution  of  ammo- 
nium carbonate.  The  voluminous  precipitate  which  first  falls  is 
acted  upon  by  an  excess  of  the  precipitant,  sometimes  dissolving 
completely,  and  crystalline  ammonium  magnesium  carbonate, 
MgCO3.(NH4)2CO3.4H2O,  is  soon  formed;  after  standing  twenty- 
four  hours  the  precipitate  is  filtered  off,  washed  with  the  con- 
centrated ammoniacal  solution  of  ammonium  carbonate,  dried 
and  strongly  ignited.  In  the  absence  of  salts  of  potassium,  the 
residue  is  weighed  at  once  as  magnesium  oxide,  and  from  the  fil- 
trate sodium  salts  are  recovered  by  evaporation.  When  a  salt  of 
potassium  is  originally  present,  with  or  without  a  salt  of  sodium, 
the  ignited  magnesium  oxide  is  to  be  washed  out  and  again 
ignited  before  weighing,  and  the  washings  are  to  be  added  to  the 
filtrate  containing  the  greater  part  of  the  alkalies. 

Gooch  and  Eddy§  have  shown  that  ammonium  magnesium 
carbonate  is  noticeably  soluble  in  Schaffgotsch's  solution  of  full 
strength,  and  rather  more  so  in  the  same  reagent  of  half  strength, 
and  that  an  exact  separation  of  magnesium  from  the  alkalies,  in 
solutions  of  reasonable  volume,  cannot  be  made  without  modi- 
fication of  the  method.  By  suitable  addition  of  alcohol,  how- 
ever, it  has  been  found  possible  to  make  the  precipitation  com- 
plete and  to  effect  the  separation  of  magnesium  from  small 
amounts  of  alkali  salt  in  one  operation ;  when  considerable  quan- 

*  Franke  Stuart  Havens  and  Arthur  Fitch  Way,  Am.  Jour.  Sci.,  [4],  viii,  217. 

t  See  page  507. 

t  Ann.  Phys.,  civ,  482  (1858). 

§  F.  A.  Gooch  and  Ernest  A.  Eddy,  Am.  Jour.  Sci.,  [4],  xxv,  444. 


MAGNESIUM 


155 


titles  of  alkali  salt  are  present  the  separation  may  be  made  by 
two  treatments. 

The  solution  containing  the  salts  of  magnesium  and  the  alkalies 
is  brought  to  a  volume  of  about  50  cm.3  and  an  equal  amount  of 
absolute  alcohol  is  added,  precipitation  is  made  by  addition  of 
50  cm.3  of  the  saturated  ammoniacal  ammonium  carbonate  solu- 
tion containing  50  per  cent  alcohol,  and  the  mixture  is  allowed 
to  stand  twenty  minutes  after  stirring  for  five  minutes.  If  the 
amount  of  alkali  salt  originally  present  is  small,  not  exceeding 
o.i  grm.,  the  precipitate  may  be  collected  on  asbestos  in  a  perfo- 
rated crucible,  washed  with  the  precipitant,  dried,  ignited,  and 
weighed  as  magnesium  oxide.  When  the  amount  of  alkali  salt 
originally  present  is  larger,  the  precipitate  may  be  freed  from 
traces  of  the  alkali  salt  by  pouring  off  the  supernatant  liquid 
through  the  prepared  asbestos  filter,  dissolving  the  precipitate, 
and  precipitating  ammonium  magnesium  carbonate  as  at  first. 
This  second  precipitate,  collected  upon  the  filter  originally  used, 
leaves  upon  ignition  practically  pure  magnesium  oxide.  The 
accompanying  figures  show  excellent  results. 

Precipitation  of  Magnesium  Ammonium  Carbonate  and  Weighing  of  the  Oxide. 


i 

i 

.2 

_i 

'o  a 

1 

NaCl  taken 

KCI  taken. 

NH4C1  tak 

MgO  weigh 

Error  MgO 

Jfc 
*l| 

gctfS 

Volume  of 
water 
solution. 

Volume  of 
alcohol 
added. 

Volume  of 
precipitant 

Volume  of  s 
tion  used  i 
washing. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

cm.3 

cm.s 

cm.3 

cm.* 

Single  precipitation. 


0.1444 

0.1443 

—  o.oooi 

o  .  oooo 

50 

50 

50 

0.1444 

o  .  1440 

—0.0004 

O.OO02 

50 

50 

50 

.  ..  • 

o  .  1444 

O.  I 

0.1445 

+O.OOOI 

O.OOOI 

50 

50 

50 

50 

0.1444 

O.I 

0.1444 

o.oooo 

O.OOOI 

50 

50 

50 

50 

0.1444 

O.I 

0.1445 

+O.OOOI 

O.OOO2 

50 

50 

50 

50 

0.1444 

O.  I 

.  .  . 

0.1449 

+0.0005 

O.OOOI 

50 

50 

5° 

50 

0.1444 

O.  2 

o  .  1449 

+0.0005 

0.0002 

50 

50 

50 

50 

0.1444 

0.2 

o.  1461 

+0.0017 

0  .  OOOO 

50 

50 

50 

50 

0.1444 

3-o 

0.1444 

o.oooo 

O.OOOI 

50 

50 

5° 

50- 

0.1444 

3-0 

0.1447 

+0.0003 

O.OOO2 

50 

50 

50 

50 

Double  precipitation. 


0.1444 
0.1444 

0.2 

O.2 

o.  1446 

0.1442 

+O  .  OOO2 
—  O.OO02 

O.OOO2 
0.0002 

50 
50 

50 
50 

50 
50 

50 

50 

156  METHODS  IN  CHEMICAL  ANALYSIS 

The  Determination  of  Magnesium  as  the  Pyrophosphate. 

As  Neubauer*  had  previously  pointed  out,  the  ideal  ammonium 
magnesium  phosphate,  NH4MgPO4,  which  yields  upon  ignition 
the  pyrophosphate,  Mg2P2O7,  may,  according  to  conditions  of 
precipitation,  be  contaminated  by  the  trimagnesic  phosphate, 
Mg3P2O8,  or  by  a  double  phosphate,  (NH4)4Mg(PO4)2,  which  upon 
ignition  leaves  magnesium  metaphosphate.  Gooch  and  Austin  f 
have  emphasized  the  effect  of  ammonium  salts  in  bringing  about 
the  contamination  of  the  precipitate  which  results  in  the  forma- 
tion of  the  metaphosphate  on  ignition,  and  have  also  pointed 
out  that  the  use  of  strongly  ammoniacal  solutions  and  strongly 
ammoniacal  wash-water  is  distinctly  disadvantageous,  as  well 
as  inconvenient.  Having  shown  that  as  little  as  o.oooi  grm.  of 
magnesium  oxide  may  be  precipitated  in  500  cm.3  of  faintly 
ammoniacal  solution  (even  when  containing  as  much  as  60  grm. 
of  ammonium  chloride  or  100  cm.3  of  a  saturated  solution  of 
ammonium  oxalate),  the  authors  recommend  the  use  of  faintly 
ammoniacal  solutions  and  wash -water,  and,  to  prevent  as  much 
as  may  be  the  formation  of  the  phosphate  containing  excess  of 
the  ammonium  ion,  restriction  of  the  amount  of  ammonium  chlo- 
ride present.  When  ammonium  salts  are  present  in  quantity, 
as  is  the  case  in  the  ordinary  course  of  analysis,  the  precipitate 
first  thrown  down  by  addition  of  ammonium  sodium  phosphate 
and  ammonia  in  faint  but  distinct  excess  is  settled,  and  the 
supernatant  liquid  is  poured  off  through  the  filter  used  sub- 
sequently in  collecting  the  precipitate.  The  precipitate  is  dis- 
solved in  the  least  possible  amount  of  hydrochloric  acid  and 
thrown  down  again  from  the  diluted  solution  by  ammonia  in 
slight  excess.  For  safety,  a  little  ammonium  sodium  phosphate 
may  also  be  added.  The  precipitate  is  filtered  off  and  washed 
with  faintly  but  distinctly  ammoniacal  water,  and,  to  avoid 
reduction,  the  ignition  is  made  slowly  and  carefully  so  that 
all  ammonia  is  expelled  before  the  temperature  is  raised  to 
redness. 

When  filtrations  are  made  on  asbestos  in  the  perforated 
crucible,  as  was  done  in  the  experiments  the  results  of  which  are 
recorded  below,  it  is  well  to  cap  the  crucible  and  moisten  the 

*  Zeit.  Angew.  Chem.,  1896,  439. 

t  F.  A.  Gooch  and  Martha  Austin,  Am.  Jour.  Sci.,  [4],  vii,  187. 


MAGNESIUM 


157 


precipitate  upon  the  felt  with  a  drop  of  a  saturated  solution  of 
ammonium  nitrate  before  proceeding  to  dry  and  ignite.  The 
accompanying  figures  show  the  accuracy  which  may  be  expected 
when  precipitations  are  made  in  presence  of  varying  amounts 
of  ammonium  salts. 

Effect  of  Ammonium  Salts  in  Respect  to  the  Constitution  of  the  Precipitate. 


Mg2P207 
correspond- 
ing to 
Mg(N03)2 

Mg2P207 
found. 

Error  in 
terms  of 
Mg2P207. 

Error  in 
terms  of 
MgO. 

NH4C1 
present. 

HNa- 
NH4P04 
4H20 

Volume. 

taken. 

I. 

II. 

I. 

II. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

cm.3 

cm.J 

0-S3II 

0.5312 

+O.OOOI 

O.OOOO 

* 

* 

2-5 

150 

IOO 

0-53U 

0-53" 

O.OOOO 

O.OOOO 

* 

* 

2-5 

150 

IOO 

0-53II 

0.5346 

+0.0035 

+0.0013 

2 

2 

2-5 

150 

IOO 

0-53II 

0.5348 

+0.0037 

+0.0014 

2 

2 

2-5 

150 

IOO 

0-53II 

0.5383 

+0.0072 

+0.0026 

5 

5 

2.5 

ISO 

IOO 

0-53II 

0.5368 

+0.0057 

+O.OO2I 

5 

5 

2.5 

150 

IOO 

0-53II 

0.5376 

+0.0065 

+0.0023 

IO 

IO 

2.5 

2OO 

IOO 

0-53II 

0-5395 

+0.0084 

+0.0030 

IO 

IO 

2.5 

2OO 

IOO 

0-53II 

0.5396 

.+0.0085 

+0.0031 

60 

5 

2.5 

250 

IOO 

0.53H 

0.5389 

+0.0078 

+0.0028 

60 

5 

2.5 

250 

IOO 

*  Probably  less  than  i  grm. 

Gooch  and  Austin  point  out  that  good  results  are  obtained 
in  one  precipitation  by  the  method  of  Wolcott  Gibbs,  proposed 
many  years  ago.*  According  to  this  method  the  boiling  solution 
of  the  magnesium  salt  is  treated  with  ammonium  sodium  phos- 
phate, and  ammonia  is  added  after  cooling.  Even  in  presence  of 
considerable  amounts  of  ammonium  chloride  this  process  yields 
a  phosphate  of  nearly  ideal  constitution  if  only  the  boiling  be 
prolonged  from  three  to  five  minutes.  The  greater  part  of  the 
ammonium  magnesium  phosphate  —  about  90  per  cent  —  forms 
in  this  process  before  free  ammonia  is  added,  and  the  ammonium 
which  enters  the  phosphate  thus  formed  is  derived  from  the  micro- 
cosmic  salt,  which  must  become  correspondingly  acidic.  Under 
these  conditions,  the  tendency  to  form  an  insoluble  ammonium 
magnesium  phosphate,  richer  in  ammonium  and  poorer  in  mag- 
nesia than  the  normal  salt,  is  slight.  In  the  process  of  Gibbs, 
as  well  as  in  the  modified  precipitation  process,  the  use  of  the 
faintly  ammoniacal  solution  and  wash-water  is  sufficient  and 
advantageous. 

*  Am.  Jour.  Sci.,  [3],  v,  114. 


158  METHODS  IN  CHEMICAL  ANALYSIS 

The  Ar senate  Process  for  the  Separation  of  Magnesium  and  the 

Alkalies. 

Browning  and  Drushel*  have  taken  advantage  of  the  insolu- 
bility of  ammonium  magnesium  arsenate  in  ammoniacal  solution, 
with  the  reducibility  of  arsenic  acid  by  hydrobromic  acid  and 
volatility  of  arsenious  bromide,  to  perfect  a  method  for  the  sepa- 
ration of  magnesia  from  the  alkali  metals,  and  the  determination 
of  these  elements.  From  a  solution  of  the  chlorides  of  magne- 
sium and  potassium  or  sodium,  the  magnesium  may  be  precipi- 
tated in  a  distinctly  but  not  strongly  ammoniacal  solution  by 
40  per  cent  to  80  per  cent  excess  of  ammonium  arsenate,  with 
brisk  stirring.  When  only  a  small  amount  of  magnesium  is 
present  in  a  relatively  large  amount  of  solution,  the  precipitate 
forms  slowly  and  becomes  complete  only  on  long  standing.  The 
precipitation  of  amounts  of  magnesium  so  small  as  not  to  be 
precipitated  immediately  by  ammonium  arsenate  may  be  brought 
about  by  freezing  the  solution,  a  process  previously  shown  to  be 
applicable  f  in  the  precipitation  of  small  amounts  of  arsenic  acid 
by  magnesia  mixture.  Precipitation  may  also  be  hastened  by 
adding  alcohol  amounting  to  15  per  cent  to  20  per  cent  of  the  mix- 
ture and  filtering  as  soon  as  the  precipitate  settles  completely. 
The  precipitate  is  collected  under  moderate  pressure  in  an  ignited 
and  weighed  perforated  crucible  containing  a  close  felt  of  fine  as- 
bestos. It  is  washed  with  40  cm.3  to  50  cm.3  of  ammoniacal  water, 
after  which  it  is  dried  at  125°  to  140°  and  carefully  ignited  and 
weighed  as  magnesium  pyroarsenate. 

It  is  shown  elsewhere  {  that  arsenic  acid  may  be  reduced  and 
volatilized  by  the  action  of  hydrochloric  acid  and  potassium 
bromide.  The  removal  of  the  arsenic  acid  from  the  alkali  salts 
is  easily  accomplished  by  the  similar  procedure  of  treating  the 
mixture  with  hydrobromic  acid  or  with  ammonium  bromide  and 
hydrochloric  acid  and  evaporating  in  an  open  dish  under  a  good 
hood.  The  complete  method  as  recommended  for  the  estimation 
of  magnesium  and  its  removal  from  the  alkalies,  and  the  sub- 
sequent estimation  of  the  alkalies,  is  as  follows : 

The  magnesium  is  precipitated  in  a  distinctly  but  not  strongly 
ammoniacal  solution  by  a  40  per  cent  to  80  per  cent  excess  of 

*  Philip  E.  Browning  and  W.  A.  Drushel,  Am.  Jour.  Sci.,  [4],  xxiii,  293. 

t  See  page  290. 

i  See  page  316. 


MAGNESIUM 


ammonium  arsenate.  The  completeness  of  the  precipitation 
may  be  hastened  by  freezing  the  solution  in  an  ice-and-salt  mix- 
ture or  by  adding  alcohol  to  about  15  per  cent  to  20  per  cent  of 
the  total  volume  of  the  solution,  which  may  vary  from  100  cm.3 
to  250  cm.3  according  to  the  amounts  of  salt  present.  The  mag- 
nesium arsenate  obtained  is  filtered  on  an  asbestos  felt  contained 
in  a  perforated  platinum  crucible,  the  crucible  and  felt  having 
been  previously  ignited  and  weighed,  and  is  dried,  ignited  and 
weighed  as  the  pyroarsenate. 
Experimental  results  follow. 

Magnesium  and  the  Alkalies. 


(NH4),AsO« 

used 
calculated 
as  As2Oa. 

grm. 

Dilu- 
tion. 

cm.3 

NaCl  or  KC1  converted  to 
Na2SO4  or  K2SO4  and  calculated 
as  Na,O  or  K2O. 

MgCl2  converted  into  Mg2As2O7 
and  calculated  as  MgO. 

Taken, 
grm. 

Found, 
grm. 

Error, 
grm. 

Taken, 
grm. 

Found, 
grm. 

Error, 
grm. 

Precipitate  stood  12  to  24  hours. 


O.  I 

IOO 

0.1194 

o.  1191 

-0.0003 

0.0199 

0.0197 

—  O.OOO2 

O.  2 

150 

0.1194 

o.  1196 

+  O.O002 

0.0399 

0.0397 

—  O.OOO2 

0-45 

250 

0.1194 

0.1195 

+  O.OOOI 

o  .  0998 

o  .  0998 

o  .  oooo 

o-45 

250 

0.1194 

0.1194 

o.oooo 

0.0998 

0.0997 

—  o.oooi 

0-45 

250 

0.2389 

0.2385 

—  0.0004 

o  .  0998 

0.0999 

+O.OOOI 

0.4 

250 

0.0478 

0.0481 

+0.0003 

o.  1198 

0.1193 

—0.0005 

o-35 

250 

0.0956 

0.0957 

+O  .  OOOI 

o  .  0998 

o  .  0996 

—  O.OOO2 

o-35 

250 

0.0956 

0.0957 

+0.0001 

o  .  0998 

0.0994 

—0.0004 

o-4S 

250 

0.0909 

0.0915 

+0.0006 

0.0998 

0.0993 

—0.0005 

O.  I 

IOO 

0-0545 

o  .  0549 

+0.0004 

0.0006 

o  .  0004 

—  O.OOO2 

Precipitation  hastened  by  alcohol. 


O.  I 

IOO 

o.  1181 

o.  1184 

+0.0003 

o  .  0040 

0.0038 

—  O.OOO2 

O.  I 

IOO 

0.1181 

0.1184 

+0.0003 

o  .  0040 

0.0038 

—  O.OO02 

O.  I 

IOO 

0.0040 

0.0038 

—  O.OOO2 

Precipitation  hastened  by  freezing. 


Oil 

IOO 

0.1181 

o.  1184 

+0.0003 

o  .  0040 

o  .  0040 

O.OOOO 

0.2 

IOO 

0.1181 

0.1183 

+O.OOO2 

0.0040 

0.0039 

—  O.OOOI 

0-45 

250 

0.1181 

0.1179, 

—  O.OOO2 

O.  IOO2 

o.  1004 

+  O.OOO2 

,  The  nitrate  is  transferred  from  the  filter  flask  to  a  platinum 
dish,  and  after  the  addition  of  10  cm.3  of  hydrochloric  acid  (sp.  gr. 
1. 20)  and  about  the  same  amount  of  hydrobromic  acid  (sp.  gr. 
1.3),  or  I  to  3  grm.  of  ammonium  bromide,  is  evaporated  to  dry- 


160  METHODS  IN  CHEMICAL  ANALYSIS 

ness  under  a  draft  hood.  The  residue  is  gently  ignited  to  re- 
move the  ammonium  salts  and  transferred  to  a  weighed  platinum 
crucible  with  a  small  amount  of  water.  A  little  sulphuric  acid 
[i :  i]  is  added,  and  the  solution  evaporated  to  remove  the  water 
and  excess  of  sulphuric  acid,  by  placing  the  crucible  on  a  triangle 
in  a  porcelain  crucible  used  as  a  radiator.  After  the  sulphuric 
acid  has  ceased  to  fume,  the  crucible  is  removed  from  the  radi- 
ator, and  after  ignition  at  the  full  heat  of  the  Bunsen  burner  the 
alkali  is  weighed  as  the  normal  sulphate. 

CALCIUM;  STRONTIUM;  BARIUM. 

The  Detection  of  Barium  and  Strontium,  Associated  with  Calcium 

and  Lead. 

In  the  ordinary  procedure  of  qualitative  analysis  the  alkali 
earth  elements  are  usually  separated  by  ammonium  carbonate, 
after  hydrogen  sulphide  and  ammonium  hydroxide  have  been 
used  to  remove  the  greater  number  of  the  bases.  It  has  long 
been  observed  that  a  considerable  part  of  the  alkali  earth,  espe- 
cially barium  and  strontium,  fails  to  appear  when  the  ammo- 
nium carbonate  is  added.  The  reasons  given  for  this  loss  have 
been  the  oxidation  of  hydrogen  sulphide  or  other  sulphides  to 
sulphates  and  the  consequent  precipitation  of  the  alkali  earth 
sulphates,  the  formation  in  alkaline  solution  of  carbonates  and 
the  consequent  precipitation  of  the  carbonates,  and  the  tendency 
of  the  large  amounts  of  ammonium  salts  which  collect  during 
the  analysis  to  interfere  with  the  precipitation  of  the  alkali  earth 
carbonates  by  ammonium  carbonate.  Various  precautions  have 
been  suggested  to  avoid  these  sources  of  error,  such  as  the  prompt 
removal  of  the  excess  of  hydrogen  sulphide  by  boiling,  the  use  of 
freshly  prepared  hydroxides  free  from  carbonate,  and  the  removal 
of  the  ammonium  salts  by  ignition  before  attempting  to  precipi- 
tate the  alkali  earth  carbonates. 

To  obviate  these  difficulties,  Browning  and  Blumenthal  *  have 
suggested  the  precipitation  of  the  insoluble  sulphates  after  the 
removal  by  hydrochloric  acid  of  mercury  in  the  mercurous  condi- 
tion, silver,  and  that  amount  of  lead  which  may  be  precipitated; 
removal  of  the  insoluble  lead  sulphate  by  treatment  of  the 

*  Philip  E.  Browning  and  Philip  L.  Blumenthal,  Am.  Jour.  Sci.,  [4],  xxxii, 
246. 


CALCIUM;    STRONTIUM;    BARIUM  l6l 

precipitate  with  ammonium  acetate;  reduction  of  the  remaining 
insoluble  sulphates  by  ignition  with  carbon;  treatment  of  the 
residue  with  acetic  acid ;  and  testing  of  the  solution  for  barium, 
strontium  and  calcium  present  as  soluble  acetates.  The  follow- 
ing method  is  suggested:  The  solution  (about  10  cm.3),  which 
may  contain  mercury  in  the  mercurous  condition,  silver,  lead, 
barium,  strontium  and  calcium,  besides  other  elements,  is  treated 
with  hydrochloric  acid  in  faint  excess  and  the  precipitated  chlo- 
rides are  filtered  off.  To  the  filtrate  are  added  about  5  grm.  of 
ammonium  acetate,  and  a  10  per  cent  solution  of  ammonium 
sulphate  to  complete  precipitation.  After  gentle  warming,  the 
alkali  earth  sulphates  are  filtered  off  and  washed  with  a  saturated 
solution  of  ammonium  acetate  until  the  washings  give  no  test 
for  lead  by  hydrogen  sulphide.  The  filtrate  and  washings  are 
reserved  for  treatment  by  the  ordinary  course  of  analysis.  To 
the  precipitated  sulphates  on  the  paper  a  small  amount  of  pure 
sugar  carbon  is  added,  the  paper  is  rolled  up,  and  the  mass  placed 
either  in  a  porcelain  crucible  with  a  cover,  or  in  a  closed  glass 
tube,  and  heated  to  full  redness  for  a  few  minutes.  The  fused 
mass  is  treated  with  about  5  cm.3  of  50  per  cent  acetic  acid  and 
warmed,  when,  if  the  alkali  earth  elements  are  present,  an  odor 
of  hydrogen  sulphide  will  generally  be  evident.  The  extract  is 
thrown  upon  a  filter  and  the  residue  washed  with  about  5  cm.3 
of  water.  The  filtrate  containing  acid  and  water  is  treated  with 
a  few  drops  of  a  solution  of  potassium  chromate  to  test  for 
barium.  The  barium  chromate  is  removed  by  filtration,  and 
the  filtrate  boiled  with  sodium  carbonate  to  precipitate  stron- 
tium and  calcium  as  the  carbonates.  If  the  precipitate  of  the 
carbonates  is  very  small,  it  may  be  dissolved  in  hydrochloric  acid 
and  tested  spectroscopically.  If,  however,  it  is  not  too  minute 
in  quantity,  it  should  be  dissolved  in  nitric  acid  after  careful 
washing,  and  the  strontium  and  calcium  separated  by  dehydra- 
tion with  amyl  alcohol. 

The  results  follow  in  the  table.  All  tests  for  strontium  and 
calcium  were  confirmed  by  the  spectroscope. 

From  these  results  it  would  appear  that  these  tests  for  barium 
arid  strontium  are  effective  to  at  least  a  milligram  of  each  ele- 
ment and  may  with  advantage  precede  the  group  precipitation 
by  hydrogen  sulphide  in  the  ordinary  course  of  qualitative 
analysis. 


162 


METHODS  IN  CHEMICAL  ANALYSIS 


Pb  present, 
grin. 

Ba  present, 
grm. 

Sr  present, 
gnu. 

Ca  present, 
grin. 

Indications. 

0.0500 

0.0500 

0.0500 

0.0500 

Good  tests  for  all. 

0.0250 

0.0250 

0.0250 

0.0250    j 

Good  tests  f  or  Pb,  Ba 
and  Sr.     Ca  faint. 

O.OIOO 

O.OIOO 

O.OIOO 

0.0100    1 

Pb  and  Ba  good.  Sr 
fair.    Ca  doubtful. 

0.0050 

0.0050 

o  .  0050 

0.0050  1 

Pb  and  Ba  good.  Sr 
fair. 

O.OOIO 

O.OOIO 

O.OOIO 

O.OOIO     < 

Pb    good.     Ba    fair. 
Sr  faint. 

O.  IOOO 

0.0050 





Pb  and  Ba  good. 

O  .  IOOO 

O.OOIO 

Pb  good.     Ba  faint. 

The  Separation  of  Barium,  Strontium  and  Calcium  by  the  Action 
of  Amyl  Alcohol  on  the  Nitrates. 

Following  in  general  known  procedure  for  the  separation  of 
sodium  and  potassium  from  lithium  by  the  use  of  amyl  alcohol,* 
Browning  t  has  developed  methods  for  the  detection  of  calcium 
and  strontium  in  association,  the  separation  and  estimation  of 
strontium  associated  with  calcium,  and  the  separation  of  barium 
associated  with  calcium  or  with  strontium. 

Experiments  with  pure  strontium  nitrate  show  that  when  the 
dry  salt  is  treated  by  dissolving  in  a  few  drops  of  water,  adding 
amyl  alcohol,  boiling  until  the  water  is  expelled  and  the  boiling 
point  rises  to  the  normal  boiling  temperature  of  the  alcohol 
(i28°-l3O°),  filtering  upon  asbestos  in  the  filtering  crucible, 
washing  with  small  amounts  of  previously  boiled  amyl  alcohol, 
and  heating  to  150°  in  an  air  bath,  nearly  the  entire  original 
amount  is  recovered.  The  quantity  which  remains  in  solution 
in  the  dehydrated  alcohol  amounts  very  regularly  to  0.0008  grm. 
of  the  nitrate  or  0.0004  grm.  of  the  oxide  for  every  10  cm.3  of 
liquid. 

When  calcium  nitrate  in  water  solution  is  similarly  treated  by 
boiling  with  amyl  alcohol,  the  salt  passes  into  solution  with  the 
exception  of  minute  portions,  not  exceeding  altogether  0.0003  grm- 
or  0.0004  grm.,  which  separate  on  the  surface  of  the  container. 
This  very  slight  residue  (apparently  the  calcium  salt  of  an  acid 
formed  by  the  action  of  nitric  acid  upon  amyl  alcohol),  when 
dried,  dissolved  in  dilute  nitric  acid,  and  again  treated  with  amyl 
*  Gooch,  Am.  Chem.  Jour.,  ix,  33. 
f  P,  E.  Browning,  Am.  Jour.  Sci.,  [3],  xliii,  50. 


CALCIUM;   STRONTIUM;    BARIUM 


I63 


Detection  of 
Strontium  and 
Calcium. 


alcohol,  separates  out  in  the  boiling,  but,  if  ignited  and  then 
dissolved  in  a  drop  of  dilute  nitric  acid,  is  not  precipitated  by 
subsequent  boiling  with  amyl  alcohol. 

To  detect  strontium  and  calcium  associated  in  the 

form  of  nitrates,  the  mixture,  not  exceeding  0.2  grm. 

(that  being  the  limit  of  the  solubility  of  calcium  nitrate 
in  5  cm.3  of  amyl  alcohol)  is  put  into  a  test  tube  and  dissolved  in 
a  few  drops  of  water,  5  cm.3  of  amyl  alcohol  are  added,  and  the 
boiling  is  carried  on  until  the  normal  boiling  point  of  the  alcohol, 
I28°-I3O°,  is  reached.  If  strontium  is  present  to  the  amount  of 
o.ooi  grm.  of  the  oxide,  a  very  decided  separation  takes  place. 
If  the  amount  is  smaller,  it  cannot  be  readily  distinguished  from 
the  residual  spots  deposited  on  the  bottom  of  the  tube  by  the 
calcium  salt.  The  alcohol  containing  the  calcium  salt  dissolved 
is  decanted  upon  a  dry  filter  paper  in  a  dry  funnel  and  the  residue 
washed  in  the  tube  with  about  5  cm.3  of  absolute  ethyl  alcohol, 
this  also  being  filtered  into  the  tube  containing  the  amyl  alcohol. 
The  filtrate  is  reserved  to  be  tested  for  calcium.  The  residue,  if 
so  small  that  it  may  be  a  calcium  deposit,  is  dried  gently,  ignited 
by  agitating  the  tube  over  a  flame,  and  dissolved  in  a  drop  of 
dilute  nitric  acid.  To  this  solution  5  cm.3  of  amyl  alcohol  are 
added  and  the  boiling  is  repeated.  Any  amount  of  strontium 
above  0.0005  §rm-  of  tne  oxide  separates  out  distinctly,  while 
the  slight  calcium  residue  does  not  reappear,  as  is  shown  in  the 
accompanying  record. 


SrO  taken, 
grm. 

'  Ca(NO3)2  taken, 
grm. 

.  Deposit  after  first 
boiling. 

Deposit  after  second 
boiling. 

O.I 

Trace. 

None. 

O.2 

Slight. 

None. 

0.2 

Slight. 

None. 

0.2 

Distinct. 

Faintest  trace. 

0.0003 

.  . 

Faint  trace. 

Faint  trace. 

0.0003 

Faint  trace. 

Faintest  trace. 

0.0005 

Distinct. 

Distinct. 

0.0005 

Distinct. 

Distinct. 

O.OOIO 

Distinct. 

Distinct. 

O.OOIO 

O.2 

Distinct. 

Distinct. 

0.0005 

O.I 

Distinct. 

Faintest  trace. 

0.0007 

O.I 

Distinct. 

Faint  trace. 

0.0008 

0.05 

Distinct. 

Distinct. 

The  test  for  calcium  is  made  upon  the  filtrate  and  washings 
after  the  first  boiling,  by  adding  to  the  clear  liquid  about  2  cm.3 


164 


METHODS  IN  CHEMICAL  ANALYSIS 


of  dilute  sulphuric  acid.  In  five  minutes  or  less,  calcium,  if 
present  to  an  amount  exceeding  o.oooi  grm.,  appears  as  a  light, 
flocky  precipitate,  different  in  character  and  easily  distinguishable 
from  the  faint  cloudiness,  gathering  to  a  minute  and  granular 
precipitate,  which  results  from  the  presence  of  a  trace  of  stron- 
tium salt  not  precipitated  in  the  boiling  process.  Following  are 
the  tests  of  this  method. 


Sr(NO3)j  taken, 
grm. 

CaO  taken, 
grm. 

Result. 

O.OOI 
O.I 
O.2 



Faint  granular  cloudiness. 
Faint  granular  cloudiness. 
Faint  granular  cloudiness. 

O.I 
O.  I 
O.I 
0.2 

O.OOI 

0.0005 

O.OOO2 
O.OOOI 
0.00005 
O.OOO5 
O.OOO2 
O.OOOI 
O.OOO2 

Decided  flocky  floating  masses. 
Decided  flocky  floating  masses. 
Decided  flocky  floating  masses. 
Plain  flocky  floating  masses. 
Faint  flocky  floating  masses. 
Decided  flocky  floating  masses. 
Decided  flocky  floating  masses. 
Plain  flocky  floating  masses. 
Decided  flocky  floating  masses. 

Strontium  and 
Calcium. 


For  the  quantitative  estimation  of  strontium  and 

Separation  and  . 

Estimation  of  calcium,*  the  dry  nitrates  of  these  elements  are 
dissolved  in  the  least  possible  amount  of  water 
contained  in  a  small  beaker  (50  cm.3  to  100  cm.3),  a 
suitable  amount  of  amyl  alcohol  (10  cm.3  to  30  cm.3)  is  added, 
the  beaker  is  heated  upon  a  wide  piece  of  asbestos  board  so 
that  inflammable  fumes  may  not  reach  the  flame  below,  and  the 
mixture  is  boiled  with  a  thermometer  inserted  until  the  normal 
boiling  point  of  the  alcohol  (128°  to  130°)  is  reached. 

From  the  precipitated  strontium  nitrate  the  alcoholic  solution 
is  decanted  through  a  weighed  filtering  crucible  and  asbestos  felt, 
and  the  residue,  dried  at  a  gentle  heat  over  a  radiator,  is  dissolved 
in  a  drop  or  two  of  dilute  nitric  acid.  Amyl  alcohol  is  again  added 
and  the  boiling  repeated,  experience  having  shown  that  the  residue 
of  the  first  boiling  is  apt  to  retain  an  appreciable  amount  of  the 
calcium  salt.  The  precipitate  of  the  second  boiling  is  filtered 
off  upon  the  felt  through  which  the  solution  had  previously  been 
decanted,  washed  with  amyl  alcohol,  dried  at  150°,  and  weighed 
as  strontium  nitrate  Sr(NO3)2.  Correction  for  the  solubility  of 
strontium  nitrate  in  amyl  alcohol  is  made  according  to  the 
*  P.  E.  Browning,  Am.  Jour.  Sci.,  [3],  xliii,  50;  and  xliv,  462. 


CALCIUM;    STRONTIUM;    BARIUM 


quantity  of  that  reagent  left  after  the  boilings, — 0.0008  grm. 
of  the  nitrate  or  0.0004  grm.  of  the  oxide  for  every  10  cm.3  of 
amyl  alcohol  decanted  or  filtered  off,  exclusive  of  that  used  in 
washing,  in  which  process  no  appreciable  amount  of  the  stron- 
tium salt  is  dissolved. 

The  calcium  in  the  filtrate  is  determined  as  the  sulphate  by 
evaporation  of  the  solution,  ignition,  treatment  with  sulphuric 
acid,  and  a  final  ignition.  From  the  apparent  amount  of  calcium 
sulphate  found  a  correction  of  0.0005  grm-  for  the  included  stron- 
tium sulphate  is  to  be  subtracted.  Results  of  experimental  tests 
are  given  in  the  tabular  statement. 

Separation  and  Estimation  of  Stronium  and  Calcium. 
Final  volumes  25  cm.3;   two  treatments. 


SrO  taken, 
grm. 

SrO  found.* 
grm. 

Error, 
grm. 

CaO  taken, 
grm. 

CaO  found,  f 
grm. 

Error, 
grm. 

0.0148 

0-0155 

+0.0007 

0.0256 

0.0254 

—  0.0002 

0.0183 

0.0183 

O.OOOO 

o.  1030 

0.1015 

—  0.0015 

0.0364 

0.0366 

+O.O002 

0.0516 

%  0.0511 

—  0.0005 

0.0365 

0.0365 

O.OOOO 

0.0515 

0.0513 

—  O.OOO2 

0.0493 

o  .  0494 

+0.0001 

0.0515 

0.0502 

—  0.0013 

0.0497 

0.0497 

0  .  0000 

0.0519 

0.0511 

—  O.OOOS 

0.0497 

0.0503 

+0.0006 

o  .  0249 

0.0245 

—  O.OOO4 

0.0729 

0.0732 

+o  .  0003 

0.0257 

0.0251 

—  O.OOO6 

0.0730 

0.0732 

+O.OO02 

0.0255 

0-0255 

O.OOOO 

0.0744 

0.0744 

O.OOOO 

0.0258 

0.0260 

+0.0002 

0.0912 

O.O9IO 

—  O.OOO2 

0.1286 

o.  1276 

—  o.ooio 

*  Corrected  by  addition  of  0.0020  grm.  for  two  treatments  in  final  volumes  of  25  cm.3, 
t  Corrected  by  subtraction  of  0.0035  grm.  for  SrSO4  included  in  CaSO4  obtained  by  evaporation 
and  ignition. 

Final  volumes  8  cm.3;   two  treatments. 


SrO  taken, 
grm. 

SrO  found.* 
grm. 

Error, 
grm. 

CaO  taken, 
grm. 

CaO  found.  f 
grm. 

Error, 
grm. 

0.0570 

0.0571 

+O.OOOI 

0-0534 

0.0536 

+0.0002 

0-0573 

0-0573 

O.OOOO 

0.0534 

0-0539 

+0.0005 

0.0285 

0.0280 

—  0.0005 

0.0272 

0.0272 

O.OOOO 

0.0568 

0.0566 

—  O.OOO2 

0-0535 

0-0533 

—  O.OOO2 

0.0568 

0.0567 

—  0.0001 

0-0533 

0.0531 

—  O.OOO2 

0.0288 

0.0286 

+O  .  OOO2 

0.0271 

0.0268 

—  0.0003 

o.  1420 

o.  1422 

+O.OOO2 

0.0535 

o  .  0540 

+0.0005 

O.I4I9 

o.  1422 

+0.0003 

0.0665 

0.0665 

O.OOOO 

O.H35 

0.1138 

—  0.0003 

o.  1066 

o.  1066 

O.OOOO 

O.II37 

0.1132 

—  0.0005 

0.1064 

0.1066 

+O.OOO2 

*  Corrected  by  addition  of  0.0006  grm.  for  two  treatments  in  final  volumes  of  8  cm.3. 

t  Corrected  for  o.ooio  grm.  of  SrSO4  included  in  CaS04  obtained  by  evaporation  and  ignition. 


i66 


METHODS  IN  CHEMICAL  ANALYSIS 


Estimation  of 
Barium  and 
Calcium. 


Procedure  similar  to  that  employed  in  the  sepa- 
ration and  estimation  of  strontium  and  calcium  by 
the  action  of  amyl  alcohol  on  the  nitrates  may  be 
applied  to  the  separation  of  barium  and  calcium,*  but,  barium 
nitrate  being  almost  entirely  insoluble  in  the  amyl  alcohol,  there 
is  in  this  case  no  advantage  in  keeping  the  volume  of  alcohol  at 
the  lowest  point.  A  convenient  volume  at  the  beginning  of  the 
dehydration  is  30  cm.3,  and  the  results  of  one  treatment  are  fully 
as  satisfactory  as  those  of  the  double  treatment.  In  separating 
barium  from  calcium  by  this  method,  the  dry  nitrates  are  treated 
in  ioo-cm.3  beakers  by  dissolving  in  a  few  drops  of  water,  adding 
30  cm.3  of  amyl  alcohol,  boiling  until  the  normal  boiling  point 
of  the  alcohol  is  reached  (128°  to  130°),  filtering  by  means  of  the 
perforated  filtering  crucible  fitted  with  the  asbestos  felt,  washing 
with  previously  boiled  amyl  alcohol,  drying  at  150°  and  weighing. 
The  test  results  show  the  exactness  of  the  method. 

Estimation  of  Barium  and  Calcium. 


BaO  taken, 
grin. 

BaO  found, 
grm. 

•         Error, 
grm. 

CaO  taken, 
grm. 

CaO  found, 
grm. 

Error, 
grm. 

o.  1410 

o.  1406 

—  0.0004 

O.OII2 

O.OII4 

+O.OOO2 

0.1300 

O.I30I 

+0.0001 

0.0926 

0.0926 

O.OOOO 

0.1043 

0.1049 

+O.OOO6 

0.0741 

0.0736 

-0.0005 

0.0781 

0.0781 

o.oooo 

0.0556 

0-0554 

—  O.OOO2 

0.0525 

0.0526 

+O.OOOI 

0.0373 

0.0372 

—  o.oooi 

Estimation  of  The  dry  nitrates  of  the  three  elements,  barium, 
Barium  and  strontium  and  calcium,  are  dissolved  in  the  least  pos- 

Strontium  .,  ,  .  -11  / 

together,  and  sible  amount  of  water,  a  suitable  amount  (15  cm.3 
of  Calcium.  £O  ^Q  cm.3)  of  amyl  alcohol  is  added,  and  the  mixture 
boiled  until  the  normal  boiling  temperature  of  the  alcohol  is 
reached.  The  alcoholic  solution  is  decanted  from  the  precipi- 
tated barium  nitrate  and  strontium  nitrate  through  a  weighed 
filtering  crucible  and  asbestos  felt.  The  residue  is  dried  over  a 
radiator,  dissolved  in  a  drop  or  two  of  dilute  nitric  acid,  and  again 
treated  as  before  with  amyl  alcohol.  The  precipitate  of  the 
second  boiling  is  filtered  off  upon  the  asbestos  felt  previously 
used,  washed  with  amyl  alcohol,  dried  at  150°  and  weighed. 
The  calcium  is  determined  as  the  sulphate  in  the  combined 

*  P.  E.  Browning,  Am.  Jour.  Sci.,  [3],  xliii,  314. 


CALCIUM;    STRONTIUM;   BARIUM 


I67 


filtrates  and  washings.  Results,  corrected  for  the  solubility  of 
the  strontium  nitrate  (0.0008  grm.  to  10  cm.3  of  alcohol  decanted 
or  filtered,  exclusive  of  washings)  and  for  the  contamination  of 
the  calcium  sulphate  by  strontium  sulphate  (0.0005  grm.  for 
IO  cm.3  of  alcohol),  are  given  in  the  table. 

Estimation  of  Barium  and  Strontium,  and  Calcium. 


Ba(NO3)2 

Ba(N03)2 

Error 

and 
Sr(N03)2 
taken. 

and 
Sr(N03)2 
found  and 
corrected. 

Error  in 
nitrates. 

averaged  and 
calculated  as 
oxide. 

CaO  taken. 

CaO  found. 

Error. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

0.3941 

0-3945 

+0  .  0004* 

+O.OOO2 

0.0283 

0.0277 

—0.0006 

0.1436 

0.1442 

+0.0006* 

+0.0003 

0.0568 

0.0558 

—  o.ooio 

0.3163 

0.3IS2 

—  O.OOII* 

—  0.0006 

0.0284 

0.0274 

—  O.OOIO 

0.1978 

0.1987 

+0.0009* 

+0  .  0005 

0.0285 

0.0280 

—0.0005 

0.1948 

0.1932 

—  o.ooi6f 

—0.0008 

0.0833 

0.0835 

+O.OOO2 

0.1971 

0.1971 

o.oooo* 

O.OOOO 

o  .  0830 

0.0817 

—0.0013 

0.1973 

0.1960 

—0.0013* 

—0.0007 

0.0830 

0.0824! 

—0.0006 

0.1959 

0.1970 

+O.OOII* 

+0.0005 

0.0830 

0.0819 

—  O.OOII  . 

0.1971 

0.1963 

—  o.oooSf 

—o  .  0004 

0.0834 

0.0831^ 

—0.0003 

*  Final  volume  in  each  of  two  treatments,  30  cm.s. 
t  Final  volume  in  each  of  two  treatments,  15  cm.3. 
J  CaS04  precipitated  and  filtered:  in  other  experiments  obtained  by  evaporation. 

The  Separation  of  Barium  and  Strontium  by  the  Action  of  Amy  I 
Alcohol  on  the  Bromides  * 

Methods  upon  which  dependence  can  be  placed  for  the  separa- 
tion of  barium  and  strontium  are  few  in  number.  The  differences 
in  solubility  of  the  bromides  of  barium  and  strontium  in  amyl 
alcohol  provide  a  method  for  a  comparatively  good  separation 
and,  when  properly  corrected,  an  exact  determination  of  these 
elements.  Anhydrous  barium  bromide  dissolves  in  amyl  alcohol 
to  the  extent  of  about  0.0013  grm.  in  10  cm.3,  while  the  same 
quantity  of  amyl  alcohol  will  take  into  solution  approximately 
0.2  grm.  of  strontium  bromide.  When  a  mixture  of  the  dry  salts 
is  dissolved  in  water,  and  the  water  removed  by  boiling  in  a 
suitable  amount  of  amyl  alcohol,  the  barium  bromide  becomes 
nearly  insoluble,  while  nearly  all  the  strontium  bromide  goes  into 
solution.  The  insoluble  barium  bromide  cannot,  however,  be  fil- 
tered off,  washed,  and  dried  to  constant  weight  without  decom- 
position, so  it  becomes  necessary  to  determine  the  barium  in 
*  Philip  E.  Browning,  Am.  Jour.  Sci.,  [3],  xliv,  459. 


i68 


METHODS  IN  CHEMICAL  ANALYSIS 


some  other  form.  Moreover,  a  single  treatment  by  boiling 
does  not  remove  the  strontium  completely  from  the  precipitated 
barium  bromide.  To  separate  barium  and  strontium  taken  as 
the  bromides,  therefore,  the  mixed  salts  are  treated  in  a  beaker  by 
dissolving  in  the  least  possible  amount  of  water,  adding  10  cm.3 
of  amyl  alcohol,  and  boiling  until  the  temperature  rises  to  the 
normal  boiling  point  of  the  alcohol  (128°  to  130°).  The  solution 
is  decanted  through  a  weighed  filtering  crucible  and  asbestos  felt. 
The  gently  dried  residue  is  dissolved  in  the  minimum  amount  of 
dilute  hydrobromic  acid  and  again  boiled  with  10  cm.3  of  amyl 
alcohol.  The  precipitate  is  filtered  off  upon  the  asbestos  felt 
previously  used  and  dissolved  from  the  felt  with  water.  From 
the  solution  the  barium  is  precipitated  as  the  sulphate,  which  is 
dried,  ignited  and  weighed.  The  strontium  is  precipitated  from 
the  united  filtrates  and  washings  by  sulphuric  acid  after  the 
addition  of  ethyl  alcohol  to  insure  thorough  mixture.  The  re- 
sults of  tests  of  the  method  are  given  below. 

Estimation  of  Barium  and  Strontium. 


BaO  taken, 
grm. 

BaO  found.* 
grm. 

Error, 
grm. 

SrO  taken, 
grm. 

SrO  found.  t 
grm. 

Error, 
grm. 

O.I2I2 

0.  1219 

+0.0007 

0.1068 

0.1071 

+0  .  0003 

o.  1215 

0.  1219 

+0.0004 

0.0358 

0-0359 

+O.OOOI 

O.I22O 

O.I22I 

+O.OOOI 

0.0353 

0.0347 

—0.0006 

O.  1212 

O.  I22O 

+0.0008 

0.0363 

0.0358 

—  0.0005 

o.  1219 

O.I22I 

+O.OOO2 

0.0361 

0.0354 

—  0.0007 

O.I2II 

0.1218 

+O.OOO7 

o.  1126 

o.  1116 

—  O.OOIO 

O.I3I9 

0.1319 

0  .  0000 

0.0577 

o  .  0586 

+0.0009 

o  .  0496 

0.0492 

—0.0004 

0.0574 

0.0579 

+0.0005 

*  Corrected  for  0.0025  BaO  corresponding  to  the  bromide  dissolved  in  two  treatments  with 
10  cm.3  of  amyl  alcohol. 

t  Corrected  by  subtraction  of  0.0040  grm.  for  the  barium  sulphate  corresponding  to  dissolved 
barium  bromide. 

The  method  is  rapid,  and,  while  the  correction  to  be  applied, 
owing  to  the  solubility  of  the  barium  salt,  is  large,  it  is  definite. 

The  Estimation  of  Barium  as  the  Sulphate. 

in  Presence  of  ^n  t^ie  ordinary  mode  of  precipitating  barium  as 
Hydrochloric  barium  sulphate,  three  conditions  are  carefully  ob- 
served, —  absence  of  excess  of  acid,  slow  mixing  of 
the  reagents,  and  standing  twelve  hours,  or  until  the  precipitate 
has  completely  subsided  before  filtration.  Usually,  in  this 


CALCIUM;    STRONTIUM;   BARIUM  169 

process,  the  precipitate  is  thrown  out  in  a  finely  divided,  milky 
condition  and  settles  very  slowly.  Mar*  has  observed,  however, 
that  the  presence  of  hydrochloric  acid  influences  the  form  in 
which  the  sulphate  is  deposited  without  affecting  the  complete- 
ness of  precipitation,  provided  a  sufficient  excess  of  sulphuric  acid 
is  also  present. 

From  a  solution  of  0.5  grm.  of  barium  chloride  in  400  cm.3  of 
water  the  precipitate  appears  immediately  upon  the  addition  of 
sulphuric  acid,  settling  slowly,  and  this  condition  prevails  also, 
even  in  hot  solutions,  when  only  one  or  two  cubic  centimeters  of 
hydrochloric  acid  have  been  previously  added.  With  10  cm.3 
to  15  cm.3  of  strong  hydrochloric  acid  in  the  solution  heated  to 
85°  or  90°,  the  precipitate  settles  clear  in  ten  or  twelve  minutes, 
and  is  in  excellent  condition  for  filtration.  When  the  solution 
contains  50  cm.3  of  the  acid,  the  precipitate  settles  clear  in  five 
minutes.  Upon  adding  the  sulphuric  acid  to  such  very  acid 
solutions,  no  precipitate  shows  for  a  moment,  but  then  it  sepa- 
rates in  beautiful  crystalline  condition  and  falls  almost  immedi- 
ately. It  can  be  safely  filtered  with  or  without  pressure  in  ten 
minutes.  In  an  instance  cited,  2  grm.  of  barium  chlqride  were 
precipitated  in  the  presence  of  30  cm.3  of  hydrochloric  acid,  the 
precipitate  was  allowed  to  settle  clear,  and  was  then  filtered  and 
washed,  the  whole  operation  being  completed  in  seven  minutes. 
This  rapid  subsidence  of  the  precipitate  is  seen  in  hot  solutions 
only,  75°  being  the  lowest  temperature  compatible  with  the  at- 
tainment of  good  results,  and  85°  to  90°  better. 

Quantitative  experiments,  quoted  below,  show  that  precipi- 
tation is  practically  complete  in  400  cm.3  of  solution  when  sul- 
phuric acid  is  added  to  the  amount  of  10  cm.3  of  the  I  :  3  dilute 
acid  (sp.  gr.  1.28  —  one  of  acid  in  a  total  volume  of  four)  in 
presence  of  hydrochloric  acid  in  amounts  up  to  150  cm.3  of  the 
concentrated  acid,  sp.  gr.  1.20.  Considerable  amounts  of  barium 
are  precipitated  at  once,  but  when  only  a  few  milligrams  are 
present  complete  formation  of  the  precipitate  requires  more  time. 
Two  or  three  hours  are  in  every  case  sufficient.  In  filtering  on 
asbestos  in  the  perforated  crucible,  as  was  done,  care  must  be 
taken  to  use  a  very  close  felt,  on  account  of  the  very  minutely 
crystalline  nature  of  the  precipitate. 

*  F.  W.  Mar,  Am.  Jour.  Sci.,  [3],  xli,  288. 


METHODS  IN  CHEMICAL  ANALYSIS 


Precipitation  in  Presence  of  Hydrochloric  Acid. 


Bad,. 
2HaO 

taken. 

grm. 

Total 
volume. 

cm.» 

HC1 

[sp.gr.  1.20). 

cm.* 

Dilute 
H2S04 
(sp.gr.  1.28). 

cm.» 

Time 
between 
precipita- 
tion and 
filtration. 

min. 

BaSO4  found, 
grm. 

£*rror. 
grm. 

0.0050 

400 

IS 

IO 

u 

(  0.0023 
(  0.0043 

—  0.0025 
—0.0005 

0.0050 

400 

15 

10 

5 

0.0031 

—  0.0017 

0.0050 

400 

15 

IO 

10 

O.0040 

—  0.0008 

O.OIOO 

400 

IS 

IO 

IO 

0.0078 

—  0.0017 

0.0100 

400 

IS 

IO 

15 

0.0085 

—  o.ooio 

O.OIOO 

400 

IS 

10 

30 

0.0083 

—  O.OOI2 

O.OIOO 

400 

IS 

10 

60 

0.0087 

—  0.0008 

.  0.0030 

400 

is 

IO 

1  20 

O.OO24 

—0.0005 

0.0050 

400 

15 

10 

150 

0.0046 

—  O.OOO2 

,    0.5014 

400 

15 

10 

10 

0.4785 

—0.0003 

0.2227 

400 

15 

IO 

IO 

O.2I22 

—0.0005 

0.5003 

400 

is 

IO 

IO 

0-4773 

—0.0005 

0.5046 

400 

15 

10 

10 

0.4814 

—0.0005 

0.5016 

400 

15 

IO 

IO 

0.4888 

—  O.OOO2 

0.5004 

400 

150 

IO 

IO 

0.4779 

0  .  0000 

0.5001 

400 

150 

IO 

IO 

0.4776 

o.oooo 

in  Presence  of  Browning*  has  investigated  with  similar  results 
Nitric  Acid  or  the  effect  of  free  nitric  acid  and  aqua  regia  (3 :  I 
Aqua  Regia.  mixture  of  hydrochloric  acid  and  nitric  acid)  upon 
the  precipitation  of  barium  as  the  sulphate  in  presence  of  an 
excess  of  sulphuric  acid.  In  a  total  volume  of  100  cm.3  contain- 
ing 10  cm.3  of  dilute  sulphuric  acid  [1:3]  the  barium  sulphate 
falls  with  an  average  loss,  after  six  hours'  standing,  of  less  than 
o.ooio  grm.  in  the  presence  of  amounts  of  nitric  acid  up  to 
25  per  cent  of  the  entire  volume.  In  aqua  regia  the  solubility 
of  barium  sulphate  is  even  less.  Following  are  experimental 
results  of  these  methods  of  treatment. 

In  this  connection,  the  effect  of  the  presence  of  a  considerable 
amount  of  free  nitric  acid  on  the  precipitation  of  barium  as  sul- 
phate in  cases  where  certain  substances  are  present  which  under 
ordinary  conditions  tend  to  hold  up  the  precipitate,  is  of  inter- 
est. Freseniusf  has  demonstrated  this  property  in  the  case  of 
ammonium  nitrate,  Scheerer  and  RubeJ  have  shown  that  meta- 

*  Philip  E.  Browning,  Am.  Jour.  Sci.,  [3],  xlv,  399. 

t  Zeit.  anal.  Chem.,  ix,  62. 

}  Jour,  prakt.  Chem.,  Ixxv,  113-116. 


CALCIUM;    STRONTIUM;   BARIUM 

Precipitation  in  Presence  of  Nitric  Acid. 


171 


BaSO* 

equivalent 

Ba(NO,)2 
taken. 

grm. 

BaS04 

found. 

grm. 

Error  in 
terms  of 
BaSO4. 

grm. 

Averages, 
grm. 

Time  between 
precipitation 
and  filtration. 

hours. 

Per  cent  by 
volume  of 
strong  HNO3. 

Total 
volume. 

cm.» 

0.3540 

0.2336 

—  0.0004 

| 

12 

5 

IOO 

0.2489 

0.2483 

—0.0006 

12 

5 

IOO 

0.2495 

0.2489 

—  O.OOO6 

v  —  o  .  OOOO 

12 

5 

IOO 

0.2492 

0.2482 

—  O.OOIO 

J 

12 

5 

IOO 

0.2486 

o  .  2483 

-0.0003 

>  —  O.OOO2 

6 

5 

IOO 

0.2490 

o  .  2490 

o.oooo 

6 

5 

IOO 

0.2555 

0.2546 

—  0.0009 

I                         A 

i 

5 

IOO 

0.2538 

0.2534 

—0.0004 

/  "~"  O  .  OOOO 

i 

5 

IOO 

0.2486 

0.2477 

—0.0009 

12 

25 

IOO 

0.2491 

0.2490 

—  o.oooi 

12 

25 

IOO 

0.2494 

o  .  2484 

—  O.OOIO 

12 

25 

100 

0.2538 

0.2535 

-0.0003 

f—  0.0008 

12 

25 

IOO 

o.  2492 

o  .  2484 

—0.0008 

12 

25 

IOO 

0.2487 

0.2471 

—  0.0016 

12 

25 

IOO 

0.3414 

0.3407 

—0.0007 

12 

25 

IOO 

0.2489 
0.2485 

0.2481 
0.2478 

—  0.0008 
—0.0007 

>  —0.0007 

6 
6 

25 
25 

IOO 
IOO 

Precipitation  in  Presence  of  Aqua  Regia. 


BaSO4 

equivalent 
to 
Ba(N03)2 
taken. 

BaSO4 
found. 

Error  in 
terms  of 
BaSO4. 

Averages. 

Time  between 
precipitation 
and  filtration. 

Per  cent  by 
volume  of 
strong  aqua 
regia. 

Total 
volume* 

grm. 

grm. 

grm. 

grm. 

hours. 

cm.1 

0-2539 

0-2534 

-0.0005 

j 

12 

5 

IOO 

0.2540 

0-2538 

—  O.OOO2 

£  —0.0002 

12 

5 

IOO 

0.2490 

o  .  2490 

O.OOOO 

) 

12 

5 

IOO 

0.2491 
0.2488 
0.3419 

0.2492 
o  .  2484 

0.3421 

+O.OOOI 

—0.0004 

+  0.0002 

>  —o.oooi 

12 

6 
6 

5 

5 
5 

IOO- 
IOO 
IOO 

0.2491 

0.2485 

—  0.0006 

1 

12 

25 

IOO 

o.  1701 
o  .  i  708 

0.1697 
0.1705 

—  O.OOO4 
-0.0003 

^-0.0003 

12 
12 

25 
25 

IOO 
100- 

0.1710 

o.  1710 

O.OOOO 

j 

12 

25 

IOO- 

0.3415 

0.3410 

—0.0005 

) 

6 

25 

IOO- 

0.3418 

0.3418 

O.OOOO 

>  —  0.0003 

6 

25 

IOO. 

0.3412 

0.3405 

—0.0007 

—0.0007 

I 

25 

100- 

phosphoric  acid  acts  similarly,  and  Spiller*  notes  the  same  gen- 
eral effect  where  alkali  citrates  are  present.     Browning  shows 
that  these  salts  cause  no  apparent  interference  with  the  precipi- 
*  Chem.  News,  viii,  280,  281. 


172 


METHODS  IN  CHEMICAL  ANALYSIS 


tation  of  barium  in  the  presence  of  nitric  acid  amounting  to  one- 
tenth  by  volume  of  the  entire  liquid.  The  barium  sulphate 
precipitated  under  such  circumstances  is,  however,  contaminated 
with  foreign  salts  present  and  must  be  purified  in  order  that  the 
amount  of  barium  actually  present  may  be  correctly  indicated. 
The  precipitate  collected  on  paper  and  ignited,  is,  therefore, 
purified  by  dissolving  it  in  sulphuric  acid  and  recrystallizing 
according  to  the  method  of  Mar,  to  be  described.*  Results  of 
this  treatment  are  given  in  the  tabular  statement. 

Purification  of  the  Precipitate. 


BaS04 

Apparent 

Impurity  present  to  the  amount 
of  5  grm. 

eauivalent 
toBa(NO3)2 
taken. 

amount  of 
BaS04 
found. 

BaSO4  after 
purification. 

Error  after 
purification  . 

Percentage 
of  strong 
HN03  by 

volume. 

grm. 

grm. 

grm. 

grm. 

Ammonium  nitrate 

o  1710 

o  1800 

O   1702 

—  O  0008 

IO 

Ammonium  nitrate 

o.  341  < 

o  .  3440 

0.3410 

—  O  0005 

IO 

Ammonium  citrate  

0.3412 

0.3442 

0.3407 

—  0.0005 

IO 

Sodium  citrate  

0.1360 

0.1730 

0.13.66 

+O.OOO6 

IO 

Metaphosphoric  acid  

0.3461 

0-35II 

0.3470 

+0.0009 

10 

Purification  of  When  barium  is  precipitated  as  the  sulphate  the 
Precipitated  Ba-  tendency  of  the  precipitate  to  include  foreign  matter, 
Tmm  Sulphate.  •£  present,  is  very  marked.  It  has  been  the  custom 
to  attempt  the  purification  of  barium  sulphate  contaminated  by 
alkali  salts  by  digesting  in  hydrochloric  acid  the  washed  precipi- 
tate. Phinneyf  has  shown,  however,  that  dilute  hydrochloric 
acid  alone  dissolves  barium  sulphate  itself,  while  mixtures  of 
hydrochloric  acid  with  enough  sulphuric  acid  to  prevent  such 
solvent  action  do  not  completely  remove  the  impurity;  and 
Mart  has  shown  that  the  presence  of  hydrochloric  acid,  even  in 
large  excess,  does  not  prevent  contamination  of  the  precipitate 
by  alkali  salts.  After  trying  ineffectually  the  purification  of  the 
Impure  barium  sulphate  by  solution  in  strong  sulphuric  acid  and 
reprecipi tation  by  water,  Mar  §  experimented  with  the  crystalli- 
zation of  barium  sulphate  from  its  solution.  The  contaminated 
precipitate  is  dissolved  in  hot  concentrated  sulphuric  acid  and 

*  This  page. 

t  J.  I.  Phinney,  Am.  Jour.  Sci.,  [3],  xlv,  468. 

t  Am.  Jour.  Sci.,  [3],  xli,  293. 

§  Loc.  cit. 


CALCIUM;    STRONTIUM;    BARIUM 


173 


recovered  from  solution  in  crystalline  form  by  evaporation  of 
the  acid.  The  crystallized  sulphate  is  then  washed  upon  a  felt 
of  asbestos  in  the  filtering  crucible,  ignited  and  weighed.  The 
evaporation  may  be  effected  over  a  radiator  or  by  means  of  a 
ring  burner;  in  either  case,  the  process  requires  several  hours. 
The  operation  may,  however,  be  completed  safely  in  a  half-hour 
by  the  aid  of  the  Hempel  evaporating  burner.  Examples  of  the 
efficiency  of  the  method  are  given  in  the  tabular  statements. 

Degree  of  Contamination  Found  in  Precipitated  Barium  Sulphate. 


BaCl2.2H2O 
taken. 

grm. 

BaSO4  found, 
grm. 

Error, 
grm. 

HC1  in  solution. 
cm.3 

Alkaline  salts 
present. 

0.5092 

0.5032 

+0.0169 

no 

KC1O3  3  grm. 

0.5027 

0.4907 

+0.0107 

IO 

KC1O3  3  grm. 

0.5026 

0.4944 

+0.0154 

IOO 

KC1      5  grm. 

0.5045 

0-4939 

+O.OI22 

10 

KC1      5  grm. 

o  .  5020 

0.4931 

+0.0137 

10 

KC1      5  grm. 

0.5013 

o  .  4849 

+O.OO6I 

10 

NaCl    5  grm. 

Slow  Evaporation  over  Radiator  or  by  Ring  Burner. 


BaCl2.2H2O  taken, 
grm. 

BaSO4  found, 
grm. 

Error, 
grm. 

0.5029 

0.4796 

—O.OOO6 

0.5008 

0.4783 

+O.OOOI 

0.5038 

0.4810 

+O.OOOI 

0.5087 

0.4861 

+0.0003 

0.5025 

0-4795 

+0.0006 

Rapid  Evaporation  by  Hempel  Burner. 


BaCl2.2H20  taken, 
grm. 

BaSO4  found, 
grm. 

Error, 
grm. 

o  .  5050 

0.4824 

+O  .  OOO2 

0.5069 

0.4838 

O.OOOO 

0.5041 

0.4825 

+0.0021 

0.5021 

0.4812 

+O.OOI8 

0.4033 

0.4801 

—  0.0005 

The  results  of  applying  this  method  to  the  purification  of  barium 
sulphate  precipitated  in  presence  of  nitric  acid  from  solutions 
containing  citrates  or  a  metaphosphate  are  given  on  page  172. 


174  METHODS  IN  CHEMICAL  ANALYSIS 

The  Estimation  of  Barium  as  the  Chloride. 
precipitation  by       It  has  long  been  known  that  barium  chloride  is 
chtoric^ckT"     insoluble  to  a  marked  degree  in  hydrochloric  acid, 
Mixture.  but  the  difficulty  of  filtering  off  the  strong  acid  and 

washing  the  precipitate  with  strong  acid  prevented  the  early  use 
of  this  characteristic  of  the  chloride  for  the  quantitative  estima- 
tion of  barium.  The  treatment  of  strong  acid  filtrates  by  means 
of  the  asbestos  felt  in  the  filtering  crucible  is  now  an  easy  matter, 
and  the  limits  of  insolubility  of  barium  chloride  in  hydrochloric 
acid  have  been  studied  by  Mar*  with  a  view  to  developing  a 
simple  method  for  the  separation  of  barium  from  calcium  and 
magnesium.  It  has  been  shown  that  barium  chloride  is  soluble 
to  an  extent  not  exceeding  one  part  in  20,000  in  pure,  concen- 
trated hydrochloric  acid,  the  solubility  increasing  very  rapidly 
with  the  diminution  in  the  strength  of  the  acid,  while  in  con- 
centrated hydrochloric  acid  containing  ether  the  solubility  falls 
to  an  amount  not  exceeding  one  part  in  about  120,000.  To 
utilize  this  fact  for  the  separation  of  barium  from  calcium  and 
magnesium,  Mar  dissolves  the  chlorides  of  the  earths  in  the  least 
possible  amount  of  boiling  water  and  precipitates  by  25  cm.3  of 
concentrated  hydrochloric  acid  with  the  addition  of  5  cm.3  of 
absolute  ether  after  cooling.  The  acid  is  added  drop  by  drop  at 
first,  as  the  precipitate  is  thus  obtained  in  a  coarse  crystalline 
condition,  filters  very  quickly,  and  is  less  liable  to  include  foreign 
matter.  After  standing  a  few  minutes  the  precipitate  is  filtered 
on  an  asbestos  felt  in  a  perforated  crucible,  washed  with  hydro- 
chloric acid  containing  about  10  per  cent  of  ether,  and  dried  at 
I5O°-2OO°.  The  method  is  accurate  and  rapid,  and  possesses 
the  further  advantage,  when  a  number  of  determinations  are  to 
be  made,  that  the  precipitate  may  be  dissolved  off  the  felt  by 
a  little  water,  and,  after  ignition,  the  crucible  and  felt  used 
again  without  reweighing.  The  felt  upon  which  a  half-dozen 
precipitates  are  thus  treated  may  not  change  by  so  much  as 
o.oooi  grm.  in  the  process.  The  fumes  of  the  strong  acid  cause 
no  inconvenience  if  the  filtration  is  performed  in  front  of  a  good 
flue. 

The  figures  of  analysis,  given  below,  indicate  the  accuracy  of 
the  process  when  applied  to  the  pure  barium  salt. 

*  F.  W.  Mar,  Am.  Jour.  Sci.,  [3],  xliii,  521. 


CALCIUM;   STRONTIUM;   BARIUM 

The  Pure  Barium  Salt. 


175 


BaCl2.2H20. 

HC1. 

Ether. 

BaCl2. 

Error. 

grrn. 

cm.8 

cm.  s 

grm. 

grm. 

0.5008 

50 

IO 

0.4267 

—0.0002 

0.5002 

5° 

IO 

0.4257 

—  0.0007 

0.4999 

5° 

IO 

0.4252 

—  0.0009 

0.4999 

50 

IO 

0.4258 

—0.0003 

0.5003 

25 

25 

0.4259 

—  0.0005 

0.5002 

25 

5 

0.4262 

—  0.0002 

0.5099 

25 

5 

0-4344 

—  0.0003 

0.5003 

25 

5 

0.4261 

-0.0003 

Following  are  figures  which  show  the  results  obtained  in  sepa- 
rating and  determining  barium  when  associated  with  magnesium 
and  with  calcium  in  mixtures  of  the  chlorides. 

Separation  of  Barium  from  Calcium  and  Magnesium. 


BaCl2.2H2O. 

Cad,. 

HC1. 

Ether. 

BaCl2. 

Error. 

grm. 

grm. 

cm.8 

cm.s 

gnu* 

grm. 

0.5001 

0-5 

50 

10 

0.4250 

—0.0013 

0.4999 

0-5 

50 

10 

0.4250 

—  o.oon 

0.5005 

0.5 

25 

25 

0.4260 

—  0.0006 

0.5002 

0.42 

25 

5 

0.4258 

—  0.0004 

0.5001 

0-5 

25 

5 

0.4255 

—0.0008 

0.5005 

0.5 

25 

5 

0.4251 

—  0.0015 

0.5001 

0.5 

25 

5 

0.4254 

—  0.0009 

0.5001 

0.5 

25 

5 

0.4258 

—  0.0005 

0.5003 

0.5 

25 

5 

0.4261 

—  0.0004 

O.  IOO2 

3-o 

25 

5 

o  .  0842 

—  O.OOI2 

O.OIO7 

3-0 

25 

5 

o  .  0080 

—  0.0005 

BaCl2.2H20. 

MgCl3.6HO2. 

HC1. 

Ether. 

BaCl2. 

Error. 

grm. 

grm. 

cm.8 

cm.J 

grm. 

grm. 

0.4999 

0-5 

25 

5 

0.4253 

—0.0007 

0.5000 

0-5 

25 

5 

0.4257 

—0.0005 

O.IOO2 

3-o 

25 

5 

o  .  0844 

—  o.ooio 

O.OIOO 

3-o 

25 

5 

0.0077 

—  0.0008 

Precipitation  by  Gooch  and  Boynton*  have  given  procedure  for  the 
Acetyi  Chloride  precipitation  of  barium  chloride  from  water  solution 
and  its  separation  from  calcium  and  magnesium  by 
the  use  of  acetyl  chloride  to  decompose  the  water  of  the  solution 
according  to  the  reaction 

CH3COC1  +  H20  =  CH3COOH  +  HC1. 
*  F.  A.  Gooch  and  C.  N.  Boynton,  Am.  Jour.  Sci.,  [4],  xxxi,  212. 


176  METHODS  IN  CHEMICAL  ANALYSIS 

Inconvenient  violence  of  the  reaction  is  moderated  by  the  addi- 
tion of  acetone  which  mixes  in  all  proportions  with  both  acetyl 
chloride  and  water,  and  by  itself  exerts  no  appreciable  solvent 
action  upon  barium  chloride. 

When  a  mixture  of  acetone  and  acetyl  chloride,  preferably 
4:1,  is  added  slowly  to  a  very  concentrated  solution  of  barium 
chloride  in  water,  the  water  is  attacked  at  once,  hydrogen  chlo- 
ride is  liberated,  and  precipitation  begins  immediately.  If  the 
temperature  is  kept  down  during  the  process  by  immersing  in 
cool  running  water  the  vessel  in  which  reaction  takes  place,  no 
more  than  a  mere  trace  of  barium  can  be  detected  by  sulphuric 
acid  in  the  residue  left  after  evaporating  the  liquid  separated 
from  the  precipitate  by  filtration  through  asbestos.  When,  how- 
ever, the  temperature  is  allowed  to  rise,  in  consequence  of  the 
heat  liberated  in  the  reaction,  an  appreciable  amount  of  barium 
may  be  found  by  sulphuric  acid  in  the  nitrate.  It  appears  that 
when  the  acetone-acetyl  chloride  mixture  [4:1]  acts  upon  the 
cooled  concentrated  water  solution  of  barium  chloride  the  pre- 
cipitate is  the  hydrous  chloride,  BaCl2.2H2O,  only  the  water  in 
excess  of  that  needed  to  form  the  hydrous  salt  being  immedi- 
ately attacked ;  that  acetyl  chloride  by  itself  produces  only  slight 
dehydration  of  this  salt  without  marked  solubility;  and  that 
prolonged  action  of  an  acetone-acetyl  chloride  mixture  [2:1] 
results  in  appreciable  dehydration  and  considerably  increased 
solubility  of  the  salt.  When  the  acetone-acetyl  chloride  mixture 
is  added  without  cooling  to  the  water  solution  of  barium  chloride 
the  heat  of  reaction  favors  dehydration  of  the  hydrous  salt,  and 
the  anhydrous  salt  may  go  into  solution  to  the  amount  of  sev- 
eral milligrams  in  10  cm.3  of  the  precipitating  mixture.  Upon 
filtering  the  mixture  and  treating  the  filtrate  with  acetone,  with 
acetyl  chloride,  or  with  the  acetone-acetyl  chloride  mixture,  the 
dissolved  anhydrous  salt  is  not  thrown  out  of  solution,  but  the 
addition  of  a  drop  of  water  is  sufficient  to  induce  immediate 
precipitation  in  the  form  of  the  hydrous  salt. 

The  best  conditions  for  the  quantitative  precipitation  of  barium 
chloride  by  the  acetone-acetyl  chloride  mixture  are  found  in  the 
use  of  minimum  amounts  of  water,  the  preservation  of  ordinarily 
low  temperature,  a  liberal  proportion  of  acetone,  and  not  too 
prolonged  digestion  of  the  precipitate  in  the  excess  of  the  pre- 
cipitant. The  salt  to  be  analyzed  is  weighed  out  into  a  small 


CALCIUM;   STRONTIUM;    BARIUM 


177 


beaker  and  dissolved  in  I  cm.3  of  water.  The  beaker  is  cooled 
by  immersion  in  a  water  bath,  preferably  supplied  with  running 
water  at  a  temperature  of  about  15°.  To  the  cooled  solution, 
constantly  shaken,  the  acetone-acetyl  chloride  mixture  is  added 
from  a  dropping  funnel  at  the  rate  of  five  drops  to  the  second. 
The  precipitate  is  filtered  off  upon  asbestos  in  a  perforated  cru- 
cible, dried,  or  ignited,  and  weighed  as  the  anhydrous  chloride, 
BaCU.  The  best  conditions  studied  for  the  handling  of  O.I  grm. 
of  hydrous  barium  chloride  are  the  solution  of  the  salt  in  I  cm.3 
of  water,  treatment  with  30  cm.3  of  the  4  : 1  mixture  of  acetone 
and  acetyl  chloride,  washing  with  acetone,  and  drying  in  the 
air  bath  at  135°  or  at  low  redness. 

Following  are  the  results  of  experimental  tests  of  the  method 
applied  to  pure  barium  chloride. 

The  Pure  Barium  Salt. 


BaCl2  taken  as 
BaCl2.2H2O. 

grm. 

BaCl2 
found. 

grm. 

Error. 

grm* 

Water  to 
dissolve 
BaCl2.2H2O. 

cm.* 

Amount  of  mixture  and 
composition  by  volume. 

To  precipitate. 

To  wash. 

0.0859 

0.0859 

o.oooof 

5  cm.3  2:1 

10  cm.3  2: 

0.0861 

0.0854 

—0.0007! 

5  cm.3  2: 

10  cm.3  2: 

0.0861 
0.0862 

0.0858 
0.0854 

—0.0003! 
—0.0008* 

5  cm.3  2: 
6  cm.3  2: 

10  cm.3  2: 
10  cm.3  2: 

0.0857 

0.0854 

—0.0003* 

6  cm.3  2: 

10  cm.3  2: 

0.0858 

o  .  0860 

+O.OOO2* 

6  cm.3  2: 

30  cm.3  4:1 

0.0860 

0.0859 

—  O.OOOI* 

6  cm.3  2: 

30  cm.3  4:1 

0.0853 

0.0850 

—0.0003* 

6  cm.    2: 

Acetone. 

0.0854 

0.0848 

—0.0006* 

6  cm.    2: 

Acetone. 

0.0852 

0.0851 

—  0.0001* 

6  cm.    2: 

Acetone. 

0.0857 

0.0856 

—  o.oooif 

6  cm.    2: 

Acetone. 

0.0852 

o  .  0845 

—0.0007! 

6  cm.    2: 

Acetone. 

0.0855 

0.0852 

—0.0003! 

6  cm.    2: 

Acetone. 

0.0862 

0.0862 

o.oooof 

30  cm.   4: 

Acetone. 

0.0868 

0.0868 

o.oooo! 

30  cm.   4: 

Acetone. 

*  Ignited  at  low  redness, 
t  Dried  at  135°  for  i$  hours. 

The  application  of  these  conditions  to  the  separa- 
tion of  barium  from  moderate  amounts  of  calcium 
and  magnesium  proves  to  be  easily  feasible.  When 
acetone  is  added  to  the  concentrated  solution  of  calcium  chloride 
or  magnesium  chloride  in  water  two  liquid  layers  are  formed,  the 


Separation  from 
Calcium  and 
Magnesium. 


i78 


METHODS  IN  CHEMICAL  ANALYSIS 


acetone  above  and  the  aqueous  layer  below;  but  the  addition  of 
a  few  drops  of  acetyl  chloride  renders  the  liquids  miscible,  while 
further  addition  causes  no  precipitation.  When  the  4: 1  mixture 
of  acetone  and  acetyl  chloride  is  added  at  the  rate  of  five  drops 
in  the  second  to  the  solution  containing  no  more  than  0.5  grm.  of 
the  calcium  and  magnesium  salts,  barium  chloride  is  precipitated 
while  calcium  chloride  and  magnesium  chloride  are  dissolved; 
but  when  the  soluble  chloride  is  present  in  the  proportion  of 

Separation  of  Barium  from  Calcium. 


BaCl2  taken  as 
BaCl2.2H2O. 

CaCl2.2H2O 
taken. 

BaCl2  found. 

Error. 

Water  used 
to  dissolve 
salts. 

Amount  of 
mixture  [4:1! 
used. 

grm. 

grm. 

grm. 

grm. 

cm.1 

cm.» 

0.0859 

O.  IOOO 

0.0859 

o.oooo* 

30 

0.0867 

o.  1040 

0.0867 

o.oooo* 

30 

0.0868 

O.  IO22 

0.0868 

0.0000* 

30 

0.0865 

O.  IO2O 

0.0865 

o.oooo* 

30 

0.0868 

O.IOI7 

o  .  0869 

+O.OOOI* 

30 

o  .  0864 

o.  1016 

0.0861 

—0.0003* 

30 

0  .  0866 

0.3025 

0.0867 

+O.OOOI* 

i 

30 

o  .  0859 

0.5025 

o  .  0859 

0.0000* 

2 

30 

0.0860 

I.OO2O 

0.0878 

+0.0018* 

3 

30 

0.0859 

I  .OO2O 

0.0855 

—  o.ooo4f 

2 

30 

0.0864 

1.0035 

0.0867 

+0.0003! 

2 

30 

*  The  precipitant  was  added  at  first  at  the  rate  of  five  drops  in  the  second, 
t  The  precipitant  was  added  at  the  rate  of  two  drops  in  the  second  at  the  outset  and  later  of  five 
drops  in  the  second. 

Separation  of  Barium  from  Magnesium. 


BaCl2  taken  as 
BaCl2.2H2O. 

MgCl2.6H20 
taken. 

BaCl2  found. 

Error. 

Water  used 
to  dissolve 
salts. 

Amount  of 
mixture  [4  :  i] 
used. 

grm. 

grm. 

grm. 

grm. 

cm.". 

cm.s 

0.0858 

O.  IOOO 

0.0857 

—  O.OOOI* 

30 

0.0869 

0.1025 

0.0870 

+O.OOOI* 

30 

0.0858 

0.1025 

0.0858 

o.oooo* 

30 

0.0862 

O.IOIO 

0.0863 

+0.0001* 

30 

0.0858 

0.1006 

0.0860 

+O.OO02* 

30 

o  .  0860 

O.  IO2O 

0.0859 

—  O.OOOI* 

30 

0.0860 

O.  IOIO 

0.0862 

+O.OOO2* 

30 

0.0865 

0.3010 

0.0867 

+0.0002* 

\ 

30 

0.0864 

0.5000 

0.0867 

+0.0003* 

2 

3® 

0.0868 

1.0015 

0.0878 

+O.OOIO* 

3 

30 

0.0853 

I.  0010 

0.0854 

+o.oooif 

3 

30 

*  The  precipitant  was  added  at  the  rate  of  five  drops  in  the  second. 

t  The  precipitant  was  added  at  first  at  the  rate  of  two  drops  in  the  second  and  later  of  five 
drops  in  the  second. 


CALCIUM;   STRONTIUM;   BARIUM 


179 


i.o  grm.  to  o.i  grm.  of  the  barium  chloride,  the  rate  of  addition 
of  the  precipitating  mixture  should  not  be  greater  than  two  drops 
in  the  second  at  the  start  in  order  to  avoid  inclusion  of  the  soluble 
salt  in  the  insoluble  barium  salt.  Even  in  such  cases  the  mixture 
may  be  added  at  the  rate  of  five  drops  in  the  second,  after  the 
greater  part  of  the  barium  is  down.  The  experimental  results 
obtained  in  the  separation  of  o.i  grm.  of  the  barium  salt  from 
0.5  grm.  of  calcium  and  magnesium  salts  are  excellent. 

When  the  4 : 1  mixture  of  acetone  and  acetyl  chloride  is  added 
to  the  concentrated  water  solution  of  o.  I  grm.  of  strontium  chlo- 
ride a  partial  precipitation  of  the  hydrous  chloride,  SrCl2.2H2O, 
takes  place. 

The  action  of  a  4 : 1  mixture  of  acetone  and  acetyl  chloride 
upon  the  concentrated  solution  of  the  chlorides  affords  easy  and 
exact  means  for  the  separation  and  estimation  of  barium  asso- 
ciated with  calcium  and  magnesium.  It  is  not  recommended 
for  the  separation  of  barium  from  strontium. 

The  Precipitation  of  Barium  Bromide  by  Ether-Hydrobromic  Acid 

Mixture. 

Thorne  *  has  shown  that  barium  bromide  dissolved  in  the  least 
possible  amount  of  water  is  completely  precipitated  by  a  mixture 
of  concentrated  hydrobromic  acid  and  ether  in  equal  parts,  and 
that  the  precipitate  may  be  obtained  of  normal  constitution, 
BaBr2,  and  weighed  as  such  if,  after  filtering  upon  asbestos  in  the 
perforated  crucible,  it  is  treated  with  ammonium  bromide,  and 
then  gradually  heated  to  250°.  Following  are  the  results  of  test 
experiments  made  in  this  manner. 

The  Pure  Barium  Salt. 


BaBr2.2H2O 
taken. 

grm. 

HBrand 
ether  [i  :  i]. 

cm.* 

BaBr2  found, 
grm. 

BaBr2  calculated, 
grm. 

Error, 
grin. 

O.2OO8 

30 

0.1793 

0.1790 

+o  .  0003 

o.  2041 

30 

0.1822 

0.1820 

+O.OOO2 

o  .  2047 

30 

o.  1821 

0.1825 

—  0.0004 

0.2171 

30 

0.1937 

0.1936 

+0.0001 

0.3101 

30 

0.2768 

0.2765 

+0.0003 

0-503S 

30 

0.4496 

0.4490 

+0.0006 

0.5015 

30 

0.4476 

0.4473 

+o  .  0003 

*  Norman  C.  Thorne,  Am.  Jour.  Sci.,  [4],  xviii,  441. 


i8o 


METHODS  IN  CHEMICAL  ANALYSIS 


Barium  is  precipitated  completely  either  by  hydrobromic  aci'd 
or  by  hydrochloric  acid*  in  mixture  with  ether,  the  precipitate 
falling  as  bromide  or  chloride  in  proportions  according  with  the 
relative  amounts  of  these  acids  present.  In  the  presence  of  a 
great  excess  of  hydrobromic  acid  the  salt  precipitated  will  be 
essentially  bromide  even  if  the  original  salt  is  taken  in  the  form 
of  the  chloride.  Salts  of  calcium  and  of  magnesium  remain  in 
solution. 

Separation  of  Barium  from  Calcium  and  Magnesium. 


BaCl2.2H2O 
taken. 

grni. 

CaCO3. 

grm. 

MgCO3. 
grm. 

HBr  and 
ether  [i:ij. 

cm.3 

BaBr2  found. 
cm.3 

Theory  as 
BaBr2. 

grm. 

Error  in 
BaBr2. 

grm. 

0.2253 

o  2088 





30 
3O 

0.2744 
O    2</?8 

0.2741 
o  2<\4o 

+0.0003 
—  O   OOO2 

o.  3273 

3O 

O   3Q71? 

o  3082 

—  o  0007 

0.3177 

3O 

0.3864 

o  386  < 

—  o  oooi 

0.5041 

0.5000 

3O 

0.6134 

0.6143 

—  o  0009 

o  .  5083 

0.5000 

30 

0.6185 

0.6191 

—  o  0006 

o  5046 

o  5000 

3O 

o  6130 

o  6136 

-j-o  0003 

0.5022 
0.5018 
O   ^OO7 

0.5000 
0.5000 

o  3000 

30 
30 

7Q 

0.6110 
0.6106 
o  6087 

0.6104 
0.6108 
o  6002 

+0.0006 

—  0.0002 

—  o  ooo  ^ 

o  ^048 

o  3000 

•?o 

o  614.4. 

o  6142 

+O   OOO2 

The  Estimation  of  Calcium,  Strontium  and  Barium,  Precipitated 

as  Oxalates. 

The  very  high  degree  of  insolubility  which  makes  possible  the 
well-known  and  exact  process  for  the  determination  of  calcium 
by  precipitation  as  oxalate  from  ammoniacal  water  solutions,  and 
weighing,  is  not  directly  applicable  to  strontium  and  barium  on 
account  of  the  greater  solubility  of  the  oxalates  of  those  elements. 
Strontium  oxalate  is  soluble  in  12,000  parts  of  water, f  while  one 
part  of  barium  oxalate  dissolves  in  less  than  3000  parts  of  cold 
water.  % 

Gravimetric  Peters§  has  shown,  however,  that  strontium  salts 

Sestr^dum°n  may  be  precipitated  by  ammonium  oxalate  with 
and  Barium.  practical  completeness  in  a  solution  containing  one- 
fifth  of  its  volume  of  85  per  cent  alcohol,  and  with  approximate 

*  Mar,  page  174. 

t  Souchay  and  Lenssen,  Ann.  Chem.,  cii,  35. 
t  Souchay  and  Lenssen,  Ann.  Chem.,  xc,  102. 
§  Am.  Jour.  Sci.,  [4],  xii,  223. 


CALCIUM;   STRONTIUM;  BARIUM 


181 


completeness  from  water  solutions  at  a  dilution  not  exceeding 
250  cm.3  in  presence  of  an  amount  of  ammonium  oxalate  sev- 
eral times  larger  than  that  required  theoretically;  and  that  ba- 
rium oxalate  is  precipitated  almost  completely  from  a  solution 
containing  one- third  of  its  volume  of  85  per  cent  alcohol.  The 
precipitates  thrown,  down  hot  and  allowed  to  stand  over  night 
may  be  filtered  off  on  asbestos  in  the  perforated  crucible,  ignited 
a  few  minutes  in  the  flame  of  a  Bunsen  burner  and  weighed  as 
carbonate,  or,  after  treatment  with  sulphuric  acid,  as  sulphate. 
The  results  thus  obtained  are  fairly  accurate,  as  shown. 

Precipitation  in  Approximately  17  per  cent  Alcohol. 


SrO  taken  as 
Sr(N03)2. 

SrO  calculated 
from  SrCO3  found. 

Difference. 

grm. 

grm. 

grm. 

O.  1  1  20 
O.  II2O 

O.III3 
o.  1116 

—  0.0007 
—  0.0004 

0-2435 

0.2425 

—  O.OOIO 

Precipitation  in  Approximately  25  per  cent  Alcohol. 


BaO  taken  as 
Ba(N03)2. 

BaO  found  as 
BaC03. 

Difference. 

grm. 

grm. 

grm. 

0.2912 
o.  2912 

0.2912 

0.2909 
0.2901 
o  .  2901 

—  0.0003 
—  O.OOII 
—  O.OOII 

Titration  of  Peters  *  has  also  shown  that  calcium,  strontium  and 

pofaL7umbper-  barium  may  be  accurately  estimated  by  titration  of 
manganate.  the  oxalates  in  hydrochloric  acid  solution  by  potas- 
sium permanganate  in  presence  of  a  manganous  salt.f 

Calcium  oxalate  is  precipitated  from  the  boiling  hot  solution 
with  ammonium  oxalate,  and  allowed  to  stand  twelve  hours. 
The  supernatant  liquid  is  decanted  upon  asbestos  in  the  filter- 
ing crucible.  The  precipitate  is  washed  two  or  three  times, 
by  decantation,  with  50  cm.3-ioo  cm.3  of  cold  water  and  brought 
on  the  felt  with  care  to  avoid  extended  washing  with  hot  water 
after  all  the  precipitant,  ammonium  oxalate,  has  been  removed. 
The  crucible  containing  the  precipitate  is  returned  to  the  beaker, 

*  Am.  Jour.  Sci.,  [4],  xii,  216. 

t  See  page  50. 


182 


METHODS  IN  CHEMICAL  ANALYSIS 


100  cm.3-2OO  cm.3  of  water  added,  together  with  5  cm.3-io  cm.3 
of  strong  hydrochloric  acid  and  0.5  grm.-i.o  grm  of  manganous 
chloride,  and  the  oxalic  acid  titrated  at  a  temperature  of  35°-45°. 
The  results  given  are  obviously  excellent,  and  show  that  calcium, 
taken  as  the  oxalate,  may  be  estimated  by  potassium  perman- 
ganate in  the  presence  of  hydrochloric  acid  and  a  manganous 

salt. 

Precipitation  from  Water  Solution. 


CaO  taken  as 
CaCl2. 

Ammonium 
oxalate. 

Volume  at 
precipitation. 

CaO  found. 

Error. 

grm. 

grm. 

cm.* 

.grm. 

grm. 

0.0656 

0-3 

IOO 

0.0657 

-f-O.OOOl 

0  .  0656 

0-3 

100 

0.0656 

o.oooo 

o  .  0656 

0-3 

'  150 

0.0658 

+O.OOO2 

0.0656 

0-3 

IOO 

0-0655 

—  o.oooi 

0.0985 

0-5 

175 

0.0981 

—0.0004 

O.I3I3 

0.6 

150 

O.I3I5 

+O.OOO2 

O.I3I3 

0.6 

2OO 

O.I3IS 

+O.OOO2 

Peters  showed  also  that  sulphuric  acid  may  be  employed  in 
place  of  hydrochloric  acid  and  manganous  chloride  when  the 
dilution  at  titration  is  sufficient. 

In  precipitating  strontium  as  the  oxalate  in  alcoholic  solution, 
ammonium  oxalate  is  added  to  the  hot  solution  with  85  per  cent 
alcohol  amounting  to  one-fifth  to  one-third  of  the  total  volume, 
the  mixture  is  allowed  to  stand  over  night,  and  the  clear  liquid 
decanted  upon  an  asbestos  filter.  The  precipitate  is  washed 
with  a  mixture  of  equal  parts  of  85  per  cent  alcohol  and  water, 
transferred  to  the  filter,  dried  in  the  filtering  crucible  over  a 
flame  to  free  it  from  alcohol,  returned  to  the  beaker  previously 

Precipitation  from  Alcoholic  Solution. 
Volume  during  titration  150  to  250  cm.3. 


SrO  taken  as 
Sr(N03)2. 

Ammonium 
oxalate. 

Volume  at 
precipita- 
tion. 

Proportion 
of  85  per 
cent  alcohol. 

Acid 
present 
during 
titration. 

SrO  found. 

Error. 

grm. 

grm. 

cm.3 

grm. 

grm. 

0.0974 
0.0974 

0-4 
0-4 

IOO 
IOO 

t 

HC1 
HC1 

0.0973 
0.0983 

—  O.OOOI 
+o  .  0009 

0.0974 

0-4 

IOO 

£ 

HC1 

0.0975 

+  O.OOOI 

0.0974 

0.8 

IOO 

I 

HC1 

0.0981 

+O.OOO7 

o.  1948 

0.4 

2OO 

\ 

HC1 

0.1943 

—  O.OOO5 

o.  1948 

0.8 

2OO 

3 

HC1 

o.  1942 

—  O.OOO6 

CALCIUM;   STRONTIUM;  BARIUM 


dried,  treated  with  5  cm.3-io  cm.3  of  hydrochloric  acid  and 
0.5  grm.-i.o  grm.  of  a  manganous  salt,  and  the  liberated  oxalic 
acid  is  titrated  by  permanganate.  The  results  obtained  by 
this  method  are  accurate. 

Precipitation  from  Water  Solution. 


SrO  taken  as 
SrClj. 

Ammonium 
oxalate. 

Volume  at 
precipitation. 

Acid  present 
during 

SrO  found.  " 

Error. 

grm. 

grm. 

cm.s 

grm. 

grm. 

0.0974 

0.8 

100 

HC1 

0.0971 

—0.0003 

0.0974 

0.8 

100 

HC1 

0.0980 

+0.0006 

0.0974 

0.8 

100 

HC1 

0.0975 

+O.OOOI 

0.0974 

0.8 

IOO 

HC1 

0.0980 

+0.0006 

0.0974 

0.8 

IOO 

HC1 

0.0973 

—  o.oooi 

0.0974 

0.8 

IOO 

HC1 

0.0978 

+0.0004 

By  precipitating  in  the  water  solution  in  presence  of  a  con- 
siderable excess  of  ammonium  oxalate  and  washing  with  small 
amounts  of  water  (30  cm.3-4O  cm.3)  applied  judiciously,  the 
loss  by  solubility  may  be  made  practically  inappreciable,  and  in 
this  case,  there  being  no  alcohol  present  to  affect  the  titration, 
the  precipitate  need  not  be  dried  before  treatment  with  perman- 
ganate. The  results  above  show  that  o.i  grm.  of  the  strontium 
salt,  calculated  as  the  oxide,  may  be  estimated  as  the  oxalate 
with  accuracy  when  precipitated  in  100  cm.3  of  water  by  a  suf- 
ficient excess  of  ammonium  oxalate. 

Precipitation  from  Water  Solution. 


SrO  taken  as 
Sr(N08)z. 

grin* 

Ammonium 
oxalate. 

grm. 

Volume  at 
precipitation. 

cm.8 

Acid  present 
during 
titration. 

SrO  found, 
grm. 

Error. 

grrn. 

0.0974 

0-5 

IOO 

H2S04 

0.0966 

—0.0008 

0.0974 

0-5 

IOO 

H2S04 

0.0985 

+O.OOII 

0.0974 

0-5 

IOO 

H2SO4 

0.0977 

+0.0003 

0.0974 

0-5 

IOO 

H2SO4 

0.0963 

—  o.oon 

0.0974 

0.8 

IOO 

H2SO4 

0.0981 

+0.0007 

0.0974 

0.8 

IOO 

H2SO4 

0.0966 

—0.0008 

0.0974 

i  .0 

IOO 

H2S04 

0.0965 

—  0.0009 

0.0974 

2.O 

IOO 

H2SO4 

o  .  0963 

—  o.oon 

0.0974 

2.O 

IOO 

H2S04 

0.0970 

—0.0004 

0.0778 

0-5 

IOO 

H2SO4 

0.0792 

+0.0014 

0.0778 

0-5 

IOO 

H2SO4 

0.0767 

—  O.OOII 

0.0778 

o-5 

IOO 

H2S04 

0.0776 

—  O.OOO2 

0.0778 

o-5 

IOO 

H2S04 

0.0776 

—  O.OOO2 

0.0974 

0.8 

250 

H2S04 

0.0973 

+  0.0001 

0.0974 

2.O 

250 

H2SO4 

0.0975 

—o.oooi 

184 


METHODS  IN  CHEMICAL  ANALYSIS 


Amounts  of  strontium  oxalate  approximately  equivalent  to 
O.I  grin,  of  strontium  oxide  may  be  successfully  precipitated 
without  alcohol  and  titrated  in  a  volume  of  200  cm.3~3OO  cm.3 
in  presence  of  sulphuric  acid,  as  shown  by  the  results  which 
are  given  in  the  preceding  table. 

To  precipitate  barium  as  the  oxalate,  ammonium  oxalate  is 
added  to  a  solution  of  a  barium  salt,  containing  30  per  cent  of  its 
volume  of  alcohol,  and  after  standing  over  night  the  precipitate 
is  filtered  on  asbestos,  washed  by  decantation  with  100  cm.3-2OO 
cm.3  of  water  containing  30  per  cent  of  its  volume  of  alcohol, 
and  dried  over  a  flame  to  insure  the  removal  of  alcohol.  The 
crucible  containing  the  precipitate  is  returned  to  the  beaker, 
also  previously  dried  over  a  flame,  100  cm.3-2OO  cm.3  of  water, 
5  cm.3— 10  cm.3  of  strong  hydrochloric  acid,  and  0.5  grm.— i.o  grm. 
of  manganous  chloride  are  added,  and  the  solution  is  titrated  at 
35°-45°  with  permanganate.  The  results  of  the  experiments 
given  show  that  barium,  either  as  the  nitrate  or  chloride,  may 
be  estimated  in  the  manner  described  with  a  fair  degree  of 
accuracy. 

Precipitation  from  Alcoholic  Solution. 


BaO  taken  as 
Ba(N03)2. 

grm. 

Ammonium 
oxalate. 

grm. 

Volume  at 
precipitation. 

cm.* 

Acid  present 
during 
titration. 

BaO  found, 
grm. 

Error, 
gnu. 

0.1165 

0.2 

TOO 

HC1 

0.1177 

+O.OOI2 

0.1165 

O.2 

100 

HC1 

0.1170 

+0.0005 

0.1165 

O.2 

IOO 

HC1 

0.1164 

—  O.OOOI 

0.1165 

0.2 

IOO 

HC1 

0.1151 

—  O.OOI4 

0.1165 

O.2 

IOO 

HC1 

0.1165 

O.OOOO 

0.1165 

O.2 

IOO 

HC1 

0.1176 

+O.OOII 

0.1165 

O.2 

IOO 

HC1 

0.1164 

—  O.OOOI 

0.2330 

0-4 

IOO 

HC1 

0.2319 

—  o.oon 

0.2330 

0.4 

IOO 

HC1 

0.2335 

+o  .  0005 

0.2330 

0.4 

IOO 

HC1 

0.2342 

+O.OOI2 

BaO  taken  as 

BaCU. 

0  .  0942 

0.4 

IOO 

HC1 

0.0952 

+0.0010 

0.0942 

0.4 

IOO 

HC1 

0.0939 

—0.0003 

0.0942 

0.4 

IOO 

HC1 

o  .  0941 

—  O.OOOI 

0.1884 

0.4 

IOO 

HC1 

0.1893 

+0.0009 

0.1884 

0.4 

IOO 

HC1 

0.1892 

+o  .  0008 

Barium  oxalate  cannot  be  successfully  titrated  in  presence  of 
sulphuric  acid  on  account  of  the  great  insolubility  of  barium 
sulphate  and  its  protecting  influence  upon  undecomposed  barium 
oxalate. 


CHAPTER  V. 
ZINC;  CADMIUM;  MERCURY. 

ZINC. 

The  Estimation  of  Zinc  as  the  Pyrophosphate. 

IN  studying  the  determination  of  zinc  by  the  method  which 
involves  precipitation  as  ammonium  zinc  phosphate  and  weigh- 
ing as  the  ignited  pyrophosphate,  Austin*  has  shown  that, 
as  in  the  case  of  the  similar  precipitation  of  manganese,  f  the 
presence  of  a  definite  excess  of  ammonium  salt  during  the  precipi- 
tation is  essential  to  the  formation  of  the  ideal  salt,  NH4ZnPO4, 
uncontaminated  by  the  tri basic  phosphate,  Zn3(PO4)2,  while  too 
much  tends  to  produce  a  double  salt  too  rich  in  ammonia. 
The  condition  of  the  ammonium  zinc  phosphate  most  nearly 
approximating  to  the  ideal  is  obtained  by  precipitating  in  pres- 
ence of  ammonium  chloride  in  large  amount.  Microcosmic 
salt  is  added  until  the  solution  (100  cm.3  to  200  cm.3)  contain- 
ing the  ammonium  salt  is  alkaline,  and  the  whole  is  heated  until 
the  mass  subsides  in  crystalline  condition.  The  amount  of 
ammonium  chloride  should  be  20  grm.  if  the  filtration  is  to  be 
made  as  soon  as  the  solution  cools.  One-half  this  amount  will 
do  if  the  liquid  stands  a  number  of  hours.  Larger  amounts 
tend  to  give  a  salt  too  rich  in  ammonia.  The  time  of  standing 
seems  to  be  a  less  important  factor  than  either  the  excess  of 
microcosmic  salt  or  the  amount  of  ammonium  chloride.  When 
the  solutions  are  made  finally  faintly  acid  with  acetic  acid, 
according  to  the  method  of  Langmuir,{  the  results  are  low. 
Following  are  the  results  obtained  by  the  method  described. 

In  a  subsequent  article  §  it  is  made  plain  that  Dakin's  ||  pro- 
posal to  wash  with  a  I  per  cent  solution  of  ammonium  phosphate, 
followed  by  alcohol,  leads  to  erroneous  results. 

*  Martha  Austin,  Am.  Jour.  Sci.,  [4],  viii,  210. 
t  See  page  483. 

t  Jour.  Am.  Chem.  Soc.,  xxi,  115. 
§  Am.  Jour.  Sci.,  [4],  xiv,  156. 
II  Chem.  News,  Ixxxii,  101;   Ixxxiii,  37. 
185 


186 


METHODS  IN  CHEMICAL  ANALYSIS 


Estimation  as  Zinc  Pyrophosphate. 


Zn2P207 
correspond- 
ing to 
ZnS04 
taken. 

Found. 

Error. 

Error  in 
terms  of 
zinc. 

Zn2P,O7 

correspond- 
ing to  Zn 
left  in  the 
filtrate. 

HNaNH<- 
PO<.4H20. 

NH4C1. 

Time  of 
standing. 

grm. 

grm. 

grm. 

grm. 

gnu. 

grm. 

grm. 

hrs. 

0-6355 

0-6335 

—  O.OO2O 

—  0.0008 

none 

4-47 

10 

16 

0-6355 

0.6381 

+0.0026 

+O.OOIO 

none 

4-47 

20 

i 

0-6355 

0.6379 

+0.0024 

+0.0009 

none 

4-47 

20 

2 

0-6355 

0.6386 

+0.0031 

+0.0012 

none 

4-47 

20 

i 

0-6355 

0.6393 

+0.0038 

+0.0014 

none 

4-47 

20 

I 

0.6367 

0-6355 

+O.OOI2 

+0.0005 

none 

4-47 

30 

16 

The  Conversion  of  Zinc  Chloride  to  Zinc  Oxide. 

Havens*  has  shown  that  zinc  chloride  may  be  quantitatively 
converted  to  zinc  oxide  by  treatment  with  nitric  acid,  evapora- 
tion of  the  excess  of  acid,  and  ignition  of  the  residue.  The 
solution  of  zinc  chloride  is  evaporated  in  porcelain,  best  with  a 
gentle  current  of  air  playing  upon  the  surface  of  the  liquid  to 
avoid  spattering,  and  treated  repeatedly  with  nitric  acid,  added 
in  small  portions  with  intermediate  evaporations.  The  residue 
is  finally  ignited  to  convert  the  nitrate  to  oxide.  Results  are 
given  in  the  table. 

Conversion  of  Chloride  to  Oxide. 


ZnO  taken  as 
chloride. 

ZnO  found. 

Error. 

grm. 

grm. 

grm. 

.0.1019 
O.IOIO 
O.  IIOO 

o.  1016 
o.  1007 
0.1095 

—0.0003 
—0.0003 
—0.0005 

The  Electrolytic  Determination  of  Zinc. 

The  deposition  of  zinc  upon  the  rotating  crucible  f  succeeds 
best,  according  to  Medway,t  when  the  zinc  salt  —  preferably  the 
sulphate  —  is  dissolved  in  50  cm.3  of  water  to  which  4  grm.  of 
potassium  oxalate  are  added.  The  presence  of  ammonium  salts 
appears  to  retard  the  complete  deposition  of  the  metal. 

*  F.  S.  Havens,  Am.  Jour.  Sci.,  [4],  vi,  45. 

f  See  Fig.  13,  page  12. 

J  H.  E.  Medway,  Am.  Jour.  Sci.,  [4],  xviii,  56. 


ZINC 


187 


Deposition  on  the  Rotating  Cathode. 


Zinc  taken, 
grm. 

Zinc  found, 
grm. 

Error, 
grm. 

Current, 
amp. 

N.  D.™. 

Time, 
min. 

0-0553 

0.0556 

+0.0003 

2-5 

8-3 

25 

0-0553 

0-0553 

o.oooo 

2-5 

8-3 

25 

0-0553 

0.0552 

—  O.OOOI 

2-5 

8-3 

25 

0.0993 

0.0995 

+O.OOO2 

2-5 

8-3 

30 

0.0993 

0.0994 

+0.0001 

2 

6.6 

25 

0.0993 

0.0991 

—  0.0002 

2 

6.6 

25 

In  determining  zinc  by  electrolysis  with  stationary  electrodes, 
it  has  been  found  that,  when  the  attempt  is  made  to  remove 
the  zinc  from  the  platinum  upon  which  it  has  been  deposited, 
a  coating  of  platinum  black  is  left,  some  of  the  zinc  having 
amalgamated  with  the  platinum.  Only  by  dissolving  the  zinc, 
heating  the  crucible  to  redness  and  finally  making  another  appli- 
cation of  acid  can  this  black  coating  be  conveniently  removed. 
In  order  to  avoid  this  formation  of  platinum  black  it  has  been 
found  necessary  to  coat  the  platinum  with  copper  and  deposit 
the  zinc  upon  this.  The  zinc  and  copper  may  then  be  easily  re- 
moved together  by  acid.  In  depositing  the  zinc  upon  a  rotating 
cathode,  however,  it  is  found  to  be  unnecessary  to  coat  the  plati- 
num with  copper,  since  the  zinc  can  be  removed  without  any 
appearance  of  platinum  black. 

In  depositing  zinc  upon  the  rotating  cathode  from  an  acetate 
solution  containing  a  salt  of  iron,  Moody*  has  found  iron  de- 
posited with  the  zinc.f 


The  Estimation  of  Zinc  by  Precipitation  as  the  Oxalate  and  Titra- 
tion  with  Potassium  Permanganate. 

Wardt  has  shown  that  zinc  may  be  accurately  estimated  by 
precipitation  as  oxalate  and  titration  with  potassium  perman- 
ganate. To  the  boiling  water  solution  of  the  zinc  salt  oxalic 
acid  is  added,  followed  by  acetic  acid  in  large  amount.  The  pre- 
cipitate is  filtered  upon  asbestos  in  the  perforated  crucible,  and 
washed  with  small  amounts  of  water.  Crucible  and  precipitate 

*  Seth  E.  Moody,  Am.  Jour.  Sci.,  [4],  xxii,  484. 

t  See  page  67. 

t  H.  L.  Ward,  Am.  Jour.  Sci.,  [4],  xxxiii,  334. 


i88 


METHODS  IN  CHEMICAL  ANALYSIS 


are  treated  with  dilute  sulphuric  acid,  the  solution  heated  to 
boiling,  and  the  free  oxalic  acid  titrated  by  permanganate. 
Results  are  given  in  the  table. 

Determination  of  Zinc. 


Zinc  taken  as 
acetate. 

Volume  at 
precipitation. 

Oxalic 
acid. 

Acetic  acid 
added. 

Zinc  found. 

Error. 

grin. 

cm.3 

grm. 

cm.3 

grm. 

grm. 

0-0055 

IOO 

2 

IOO 

0.0056 

+O.OOOI 

0.0274 

IOO 

2 

IOO 

0.0276 

+0.0002 

0.0548 

50 

2 

50 

0-0553 

+o  .  0005 

0.0548 

IOO 

2 

IOO 

0.0550 

-f-O  .  OOO2 

0.1370 

IOO 

2 

IOO 

0.1372 

+O.OOO2 

CADMIUM. 
The  Estimation  of  Cadmium  as  the  Oxide. 

Precipitation  as  Various  objections  have  been  made  to  that  method 
Carbonate.  for  estimating  cadmium  which  involves  precipita- 
tion as  carbonate,  ignition,  and  weighing  as  oxide.  Browning 
and  Jones  *  have  shown  that  when  the  carbonate  is  filtered  upon 
an  asbestos  felt  in  a  perforated  crucible  previously  ignited  and 
weighed,  danger  of  reduction  is  obviated  and  the  process  is  sim- 
plified and  placed  in  the  category  of  good  analytical  methods. 
To  the  solution  of  the  cadmium  salt  in  about  300  cm.3  of  hot 
water  is  added  a  10  per  cent  solution  of  potassium  carbonate, 
drop  by  drop  and  with  constant  stirring.  The  solution,  with 
the  precipitate  in  suspension,  is  boiled  for  about  fifteen  min- 
utes. The  precipitated  cadmium  carbonate  becomes  granular 
and  settles  quickly,  and  is  then  filtered  upon  asbestos  in  the  per- 
forated crucible,  ignited  and  weighed.  Results  obtained  by  this 
method,  and  quoted  in  A  of  the  following  table,  show  a  slight 
plus  error  due,  as  was  shown  experimentally,  to  inclusion  of 
alkali  salt,  but  indicate  that  the  carbonate  method  can  be  suc- 
cessfully and  simply  applied  to  the  quantitative  estimation  of 
cadmium. 

The  work  of  Flora  f  fully  substantiates  these  results,  giving 
by  similar  procedure  the  analytical  figures  in  B. 

*  Philip  E.  Browning  and  L.  C.  Jones,  Am.  Jour.  Sci.,  [4],  ii,  269. 
t  Charles  P.  Flora,  Am.  Jour.  Sci.,  [4],  xx,  456. 


CADMIUM 


189 


Precipitation  as  Cadmium  Carbonate. 


CdO  taken. 

CdO  found. 

Error. 

grm. 

grm. 

grm. 

A. 

o.  1140 

0.1143 

+0.0003 

0.1142 

0.1137 

—  0.0005 

o.  1141 

0.1148 

+O.OOO7 

0.1141 

o.  1148 

+O.OOO7 

o.  1142 

o.  1146 

+o  .  0004 

0.1143 

O.II47 

+o  .  0004 

0.1143 

O.II44 

+0.0001 

0.1139 

o.  1146 

+0.0007 

o.  1270 

o.  1272 

+0.0002 

0.1279 

0.1283 

+o  .  0004 

0.1272 

o.  1281 

+0.0009 

o.  1278 

o.  1281 

+o  .  0003 

0.2556 

o.  2561 

+0.0005 

0.2550 

0.2547 

—0.0003 

0.1272 

0.1279 

+0.0007 

o.  1281 

0.1288 

+0.0007 

0.1274 

0.1278 

+o  .  0004 

0.1284 

o.  1290 

+o  .  0006 

0.1271 

0.1277 

+0.0006 

o.  1278 

o.  1285 

+0.0007 

0.2555 

0.2555 

o.oooo 

B. 

0.1277 

0.1275 

—  O.OOO2 

0.1277 

o.  1280 

+0.0003 

0.1277 

o.  1272 

—0.0005 

0.1399 

0.1391 

—0.0008 

0.1399 

0.1399 

o.oooo 

0.1703 

o.  1700 

—0.0003 

0.1703 

o  .  i  700 

—0.0003 

0.2129 

0.2128 

—  O.OOOI 

0.2129 

0.2128 

—  o.oooi 

0.2554 

0.2554 

o.oooo 

Precipitation  as  Flora*  has  also  tested  the  similar  method  by 
Hydroxide.  which  cadmium  is  precipitated  as  hydroxide  and 
weighed  as  oxide.  To  the  boiling  solution  of  the  cadmium  salt, 
about  300  cm.3  in  volume,  a  10  per  cent  solution  of  potassium 
hydroxide  is  added  drop  by  drop.  After  boiling  about  fifteen 
minutes  the  precipitate  settles  quickly  in  a  semigranular  state 
and  is  filtered  on  a  weighed  asbestos  felt  in  the  perforated  cru- 
cible, washed,  ignited  and  weighed.  The  results  are  lower  than 
those  of  the  carbonate  method. 

*  Loc.  cit. 


METHODS  IN  CHEMICAL  ANALYSIS 


Precipitation  as  Cadmium  Hydroxide. 


CdO  taken, 
gnu. 

CdO  found, 
gnu. 

Error, 
grm. 

0.1277 

0.1277 

o.oooo 

0.1277 

o.  1270 

—0.0007 

0.1277 

o  .  i  260 

—0.0017 

0.1277 

0.1286 

+o  .  0009 

0.1362 

0.1350 

—  O.OOI2 

0.1399 

0.1389 

—  O.OOIO 

0.1703 

0.1697 

—0.0006 

0.1703 

0.1693 

—  O.OOIO 

0.1703 

o.  1699 

—0.0004 

0.1788 

o.  1802 

+0.0014 

o.  2129 

0.2139 

+O.OOIO 

0.2129 

0.2128 

—  o.oooi 

While  these  figures  show  that  fair  results  may  be  obtained,  the 
hydroxide  method  is  not  comparable  with  the  carbonate  method 
as  to  accuracy  or  convenience.  The  precipitate  does  not  take 
the  same  granular  form;  it  is  hard  to  filter,  difficult  to  wash,  and 
is  removed  with  difficulty  from  the  beaker  in  which  precipitation 
takes  place. 


The  Estimation  of  Cadmium  as  the  Pyrophosphate. 

It  has  been  shown  by  Austin  *  that  cadmium  may  be  estimated 
with  accuracy  as  the  pyrophosphate.  The  precipitate  obtained 
by  making  alkaline  with  microcosmic  salt  the  nearly  neutral 
solution,  containing  ammonium  chloride  in  the  proportion  of  ten 
grams  to  one  hundred  cubic  centimeters,  is  allowed  to  stand 
several  hours,  then  filtered  off  on  asbestos  in  the  filtering  crucible, 
dried,  ignited  and  weighed.  The  cadmium  separates  out  from 
the  solution  as  a  beautiful  crystalline  mass  of  cadmium  ammo- 
nium phosphate  of  ideal  constitution.  The  conditions  must, 
however,  be  preserved  with  care;  there  must  be  no  excess  of 
ammonia,  no  free  acid,  and  no  excess  of  ammonium  salt  beyond 
the  quantity  indicated,  while  that  amount  is  necessary. 

A  criticism  of  this  method  by  Miller  and  Pagef  was  shown  in 
a  later  article  %  to  be  without  foundation. 

*  Martha  Austin,  Am.  Jour.  Sci.,  [4],  viii,  214. 
t  School  of  Mines  Quarterly,  xxii,  391. 
I  Am.  Jour.  Sci.,  [4],  xiv,  156. 


CADMIUM 


Estimation  as  Cadmium  Pyro phosphate. 


Cd2p2o7 

correspond- 
ing to 
CdCl2 

Found. 

Error. 

Error  in 
terms  of 
cadmium. 

Cd2P207 
correspond- 
ing to  Cd 
found  in 

HNaNH4- 
P04.4H2O. 

NH4C1. 

Time  of 
standing. 

taken. 

the  filtrate. 

grin. 

grm. 

grin. 

grm. 

grm. 

grm. 

grm. 

hrs. 

0.6972 

0.6976 

+o  .  0004 

+O.OOO2 

trace 

4-5 

IO 

16 

0.6972 

o  .  6969 

-0.0003 

—  O.OO02 

trace 

4-5 

10 

18 

0.6972 

0.6962 

—  o.ooio 

—  0.0006 

trace 

4-5 

IO 

16 

The  Electrolytic  Determination  of  Cadmium. 
The  deposition  of  cadmium  as  the  metal,  upon  the  rotating 
cathode,*  has  been  studied  by  Medwayf  for  the  sulphate  solu- 
tion, and  by  Flora {  for  solutions  containing  sulphuric  acid, 
acetates,  cyanides,  pyrophosphates,  phosphates,  oxalates,  for- 
mates, tartrates,  free  nitric  acid,  urea,  formaldehyde  or  acetal- 
dehyde.§ 

Deposition  from  Cadmium  taken  as  the  sulphate  to  the  amount  of 
the  Sulphuric  o.2  grm.  approximately  and  dissolved  in  50  cm.3  of 
lon'  water  containing  10  drops  of  dilute  sulphuric  acid 
may  be  successfully  deposited  upon  the  crucible  rotating  at  the 
rate  of  650—700  revolutions  per  minute.  To  avoid  solvent  action 
after  stopping  the  current,  dilute  ammonia  is  added  drop  by 
drop  to  faint  alkalinity,  while  the  current  is  still  passing  and 
after  complete  deposition  of  the  metal.  That  this  procedure  is 
satisfactory  the  following  results  of  Medwayll  show. 

Deposition  from  the  Solution  of  the  Sulphate. 


Cadmium 
taken. 

grm. 

Cadmium 
found. 

grm. 

Error, 
grm. 

Current, 
amp. 

N.  D.100. 

Time, 
min. 

0.1088 

o.  1083 

—0.0005 

2 

6.6 

IS 

0.1088 

0.1085 

—  0.0003 

2 

6.6 

15 

0.1088 

o.  1092 

+0  .  0004 

i-5 

5 

15 

0.1088 

0.1090 

+0  .  0002 

2 

6.6 

15 

0.1088 

0.1093 

+0.0005 

i.  5 

5 

12 

0.1088 

0.1093 

+o  .  0005 

2 

6.6 

10 

0.1088 

0.1087 

—  O.OOOI 

2 

6.6 

IO 

*  See  Fig.  13,  page  12. 

f  H.  E.  Medway,  Am.  Jour.  Sci.,  [4],  xviii,  56. 

J  Charles  P.  Flora,  Am.  Jour.  Sci.,  [4],  xx,  268  et  seq.,  and  292  et  seq. 

§  For  descriptions  and  results,  see  pages  191  to  195. 

II  Am.  Jour.  Sci.,  [4],  xviii,  56. 


192  METHODS  IN  CHEMICAL  ANALYSIS 

In  a  detailed  study  of  the  use  of  the  rotating  cathode  for  the 
estimation  of  cadmium,  Flora  *  calls  attention  to  the  fact  that 
the  dilution  of  the  solution  submitted  to  electrolysis  is  of  great 
importance.  It  is  advisable,  in  order  to  avoid  mechanical  loss, 
to  deposit  not  more  than  0.25  grm.  of  the  metal  upon  the  cathode, 
while  even  smaller  quantities  are  to  be  preferred.  The  current 
density  must  also  be  kept  within  limits ;  for  otherwise  a  spongy 
deposit  may  result.  Cadmium  seems  to  be  especially  apt  to 
form  these  spongy,  unweighable  deposits,  and  the  greatest  diffi- 
culties come  from  this  behavior  of  the  metal.  The  best  con- 
ditions for  the  deposition  in  presence  of  sulphuric  acid  may  be 
briefly  summarized  as  follows:  Cadmium  sulphate,  equivalent 
to  not  more  than  0.25  grm.  of  the  metal,  is  dissolved  in  45  cm.* 
to  50  cm.3  of  water ;  ten  to  fifteen  drops  of  dilute  sulphuric  acid 
are  added ;  and  the  solution  subjected  to  electrolysis  at  a  normal 
current  density  (N.  D.ioo)  ranging  between  3  amp.  and  9  amp. 
per  100  cm.2  of  surface.  It  is  not  necessary  to  heat  the  liquid, 
as  the  passage  of  the  current  soon  heats  it  sufficiently.  When 
electrolysis  is  complete,  the  excess  of  sulphuric  acid  may  be  de- 
stroyed with  a  slight  excess  of  ammonia  water,  the  current  broken, 
and  the  cathode  removed,  thoroughly  rinsed  with  water  and  alco- 
hol, and  dried  by  waving  over  a  free  flame.  If  the  deposit  is  not 
spongy  the  drying  is  a  matter  of  only  a  few  moments,  and  there 
is  no  danger  of  oxidizing  the  metallic  deposit.  If  it  is  preferred, 
the  current  may  be  reduced  by  interposed  resistance,  the  rotation 
stopped,  and  the  liquid  readily  siphoned  without  danger  of  in- 
juring the  metallic  coating. 

This  process  is  also  available  when  the  cadmium  is  taken 
as  the  chloride  if  the  volume  of  the  solution  does  not  exceed 
45  crnAf 

Deposition  from  *n  a  vomme  °f  6o  cm-3  to  65  cm-3>  containing  0.5 
Solutions  con-  grm.  to  2  grm.  of  sodium  acetate  and  a  small  amount 
taining  Acetates.  of  potassium  sulphate  to  so  regulate  the  conductivity 
that  the  normal  current  density  shall  not  much  exceed  3  amp., 
the  deposition  of  not  more  than  0.15  grm.  of  cadmium,  taken  as 
the  sulphate,  proceeds  rapidly  and  satisfactorily.  The  deposit 
under  the  conditions  is  rather  crystalline,  fairly  compact,  and 
easily  washed,  so  that  the  method  forms  one  of  the  very  best 

*  Charles  P.  Flora,  Am.  Jour.  Sci.,  [4],  xx,  268-276,  392-396,  454~455- 
t  Flora,  loc.  cit.,  page  392. 


CADMIUM 


193 


where  the  cadmium  is  taken  in  the  form  of  the  sulphate.  At 
greater  concentrations  the  precipitate  shows  a  tendency  to  spongi- 
ness,  and  it  fails  absolutely  when  the  cadmium  is  introduced  as 
the  chloride.* 

Deposition  from  Solution  Containing  Alkali  Acetate. 


Cadmium 
taken. 

NaOC2H3O. 

K2S04. 

Current. 

N.  D.™. 

E.M.P. 

Time. 

Cadmium 
found. 

Error. 

grm. 

grm. 

grm. 

amp. 

amp. 

volts. 

min. 

grm. 

grm. 

0.1118 

0-5 

I  .O 

1.0 

3-0 

8.0 

2O 

O.  II2I 

+0.0003 

0.1491 

i-5 

0-5 

0.9 

2.7 

8.0 

15 

0.1494 

+0.0003 

0.1491  . 

i-5 

o-5 

0.9 

2.7 

8.0 

15 

o  .  1496 

+o  .  0005 

Cadmium 
taken. 

NaOH. 

K2S04. 

Current. 

N.D.100. 

E.M.F. 

Time. 

Cadmium 
found. 

Error. 

grm. 

grm. 

grm. 

amp. 

amp. 

volts. 

min. 

grm. 

grm. 

O.  1491 

0-5 

1-25 

3-75 

8.0 

IO 

0.1496 

+0.0005 

0.1491 

0'.2* 

0-5 

0.8 

2.4 

8.0 

15 

0.1491 

O  .  OOOO 

0.1491 

0.2* 

0-5 

0.8 

2.4 

8.0 

15 

0.1493 

+  O.OO02 

0.1223 

o-S* 

O.2 

i  .0 

3-o 

12.  O 

20 

0.1223 

O  .  OOOO 

0.1223 

0-5* 

O.  2 

i  .0 

3-o 

12.0 

20 

0.1223 

0.0000 

*  Neutralized  by  a  slight  excess  of  acetic  acid. 


Deposition  from  deposition  of  cadmium  from  a  solution  of  the 

solutions  con-  double  cyanide  has  proved  to  be  very  satisfactory, 
les'  and  the  results  with  the  rotating  cathode  are  in  com- 
plete accordance  with  previous  work  on  this  method.  The  range 
of  conditions  of  current  and  quantity  of  electrolyte  is  broad; 
the  deposit  is  a  beautiful  silvery  plate  which  dries  very  quickly, 
and  is  so  compact  as  to  be  rubbed  off  only  with  difficulty;  and 
although  the  complete  deposition  of  the  metal  is  slower  than 
it  is  from  solutions  containing  sulphates  or  acetates,  it  is  suffi- 
ciently rapid.  Care  should  be  taken  to  avoid  foaming  of  the 
solution,  as  this  retards  somewhat  the  deposition  of  the  final 
traces  of  cadmium.  Generally,  a  volume  of  65  cm.3  to  70  cm.3 
is  found  most  satisfactory.  In  experiments  to  test  the  method 
the  cadmium  sulphate  was  treated  with  sodium  hydroxide  and 
the  precipitate  was  dissolved  in  potassium  cyanide.  The  follow- 
ing results  were  obtained.  f 


*  Flora,  loc.  cit.,  page  392. 
t  Ibid.,  page  272. 


IQ4  METHODS  IN  CHEMICAL  ANALYSIS 

•  Deposition  from  Solution  Containing  Alkali  Cyanide. 


Cadmium 
taken. 

NaOH. 

KCN. 

Current. 

N.  D.1M. 

E.M.P. 

Time. 

Cadmium 
found. 

Error. 

grm. 

grm. 

grm. 

amp. 

amp. 

volts. 

min. 

grm. 

grni. 

0.1491 

i-5 

o-5 

2-5 

7-5 

8 

35 

0.1498 

+0  .  0007 

0.1491 

i  .0 

o-S 

2.5-4.5 

7.5-I3-5 

8 

30 

o.  1490 

—  O.OOOI 

0.1223 

i-5 

I.O 

2-5 

7-5 

8 

35 

0.1225 

+O.OOO2 

Satisfactory  results  are  also  found  by  this  method  when  the 
cadmium  is  taken  as  the  nitrate  or  as  the  chloride  *  the  best  dilu- 
tion being  60  cm.3  to  65  cm.3  The  time  required  is  a  trifle  longer 
than  in  the  estimation  of  cadmium  sulphate  by  this  method,  and 
with  the  current  density  necessary  to  hasten  the  deposition  a 
considerable  tendency  to  foam  is  manifest. 

Deposition  from  Brand  f  has  recommended  the  use  of  a  solution 
Solutions  Con-  containing  sodium  pyrophosphate  for  the  electro- 

taining  Pyro-  .      KJ 

phosphates  or  lytic  estimation  of  metals,  among  others,  cadmium; 
OrthoPhosPhates.ancj  tne  fitness  of  this  solution  for  use  with  the  rotat- 
ing cathode  has  also  been  studied  by  Flora. | 

When  the  cadmium,  taken  as  sulphate  or  chloride,  is  precipi- 
tated by  sodium  pyrophosphate,  the  precipitate  dissolved  in  an 
excess  of  ammonium  hydroxide,  phosphoric  acid,  sulphuric  acid, 
or  hydrochloric  acid,  and  the  solution  submitted  to  electrolysis 
at  a  volume  of  60  cm.3,  fairly  accurate  results  may  be  obtained ; 
but  neither  is  the  method  so  accurate  as  those  previously  de- 
scribed, nor  are  the  conditions  so  flexible.  Practically  the  same 
thing  may  be  said  of  the  use  of  the  rotating  cathode  in  the  elec- 
trolysis of  cadmium  orthophosphate  dissolved  in  phosphoric  acid 
according  to  the  method  recommended  by  Smith. §  Flora  ||  has 
also  studied  the  behavior  of  solutions  containing  oxalates,  for- 
mates, tartrates,  urea,  formaldehyde  and  acetaldehyde,  nitrates 
and  free  nitric  acid.  The  results  of  this  study  and  of  the  work 
previously  described  may  be  summarized  as  follows: 

Summary.  —  Cadmium  taken  in  the  form  of  the  sulphate  may 
be  very  accurately  and  satisfactorily  estimated  by  deposition 

*  Flora,  loc.  cit.,  page  393,  and  page  454. 

t  Zeit.  anal.  Chem.,  xxviii,  581. 

J  Loc.  cit.,  page  273. 

§  Am.  Chem.  Jour.,  xii,  329. 

II  Loc.  cit. 


MERCURY  195 

upon  the  rotating  cathode  from  solutions  containing  sulphuric 
acid,  sodium  acetate  and  acetic  acid,  or  potassium  cyanide;  but 
little  less  satisfactorily  from  solutions  containing  urea,  formal- 
dehyde or  acetaldehyde ;  and  also,  with  proper  precautions,  from 
solutions  containing  pyrophosphates,  phosphates,  tartaric  acid 
or  formic  acid.  From  solutions  containing  oxalates  or  oxalic 
acid,  ammonium  tartrate  or  potassium  formate  satisfactory  de- 
posits are  not  obtained. 

When  taken  as  the  chloride,  cadmium  does  not  permit  such  a 
wide  range  of  conditions.  Nevertheless,  from  solutions  of  the 
chloride  containing  sulphuric  acid  or  potassium  cyanide,  or  the 
pyrophosphates,  the  metal  is  deposited  in  a  form  comparable 
with  that  obtained  when  cadmium  sulphate  is  taken.  Solutions 
of  the  chloride  of  cadmium  to  which  is  added  hydrogen  disodic 
phosphate  give  less  desirable  results;  while  solutions  containing 
urea,  formaldehyde  or  acetaldehyde  give  deposits  free  from  spon- 
giness  only  after  careful  regulation  of  the  conditions.  In  solu- 
tions containing  the  oxalates,  oxalic  acid,  the  formates  and  the 
tartrates,  acetates,  formic  acid  and  tartaric  acid  the  results 
were  negative.  Cadmium  nitrate  is  in  general  ill-fitted  for  elec- 
trolytic estimation,  the  cyanide  solution  being  the  only  one  from 
which  satisfactory  results  were  obtained.  From  solutions  con- 
taining one  per  cent  or  more  of  free  nitric  acid,  the  cadmium  is 
not  deposited  by  the  current. 

MERCURY. 
The  Gravimetric  Determination  of  Mercury  as  Mercurous  Oxalate. 

Mercury  taken  in  the  form  of  mercurous  nitrate  may  be  esti- 
mated as  mercurous  oxalate  precipitated  by  ammonium  oxalate 
and  dried  over  sulphuric  acid,  as  has  been  shown  by  Peters.*  It 
is  necessary,  however,  to  control  the  acidity,  dilution,  and  pres- 
ence of  mercuric  salts.  It  appears  that  5  cm.3  of  dilute  nitric 
acid,  sp.  gr.  1.15,  may  be  present  in  a  volume  of  100  cm.3,  and  that 
5  cm.3  of  the  acid  will  prevent  precipitation  of  small  amounts  of 
mercuric  salt, — o.oioo  grm.  to  0.0200  grm.,  calculated  as  mer- 
cury, depending  upon  the  amount  of  ammonium  oxalate  present 
in  excess.  According  to  the  procedure  recommended  by  Peters, 
mercurous  nitrate  dissolved  in  100  cm.3  of  water  containing  2  per 
*  C.  A.  Peters,  Am.  Jour.  Sci.,  [4],  ix,  405. 


196 


METHODS  IN  CHEMICAL  ANALYSIS 


cent  to  5  per  cent  of  dilute  nitric  acid,  sp.  gr.  1.15,  is  precipitated 
by  the  addition  of  ammonium  oxalate  in  slight  excess  with  stir- 
ring. It  is  an  easy  matter  to  keep  the  excess  of  the  precipitant 
within  the  limits  of  I  cm.3  to  2  cm.3  of  the  n/io  solution,  because 
the  mercurous  oxalate,  when  properly  stirred,  settles  rapidly. 
The  precipitate  is  collected  on  asbestos  in  a  perforated  crucible, 
washed  two  or  three  times  with  cold  water,  and  dried  to  constant 
weight  over  sulphuric  acid  at  the  ordinary  temperature,  since 
mercurous  oxalate  is  slowly  decomposed  at  temperatures  in  the 
vicinity  of  100°.  Mercury  in  the  mercuric  form  may  be  safely 
present  to  the  amount  of  10  per  cent  of  the  mercurous  salt. 

Results  of  experiments  according  to  this  procedure,  in  which 
the  drying  was  effected  by  exposure  for  fifteen  hours  or  less  at 
ordinary  temperatures  over  sulphuric  acid,  are  given  in  the 
accompanying  table. 

Mercurous  Oxalate  Dried  over  Sulphuric  Acid. 


Hg  taken  as 
Hg,(NO,),. 

Hg  present  as 
Hg(NO,),. 

Excess  of  H/IO 
ammonium 
oxalate. 

HNO3 
(sp.gr.  1.  15.) 

Volume  at 
precipita- 
tion. 

Hg  found. 

Error. 

grm. 

grm. 

cm.3 

cm.8 

cm.s 

grm. 

grm. 

o.  1217 

2-4 

IOO 

0.1217 

O.OOOO 

O.I2I7 



2-4 

IOO 

o.  1217 

O.OOOO 

O.  1122 

2-4 

IOO 

o.  1124 

+O.OOO2 

O.II22 

0.0067 

0-93 

2 

IOO 

0.1130 

+0.0008 

O.II22 

0.0067 

0-93 

2 

IOO 

O.  III2 

—  O.OOIO 

O.II22 

0.0067 

4.40 

2 

IOO 

o.  1124 

+O.OOO2 

O.II22 

0.0135 

0.72 

4 

IOO 

o.  1125 

+o  .  0003 

0.2244 

0.0071 

1.68 

4 

IOO 

0.2253 

+o  .  0009 

O.2244 

0.0071 

2.46 

4 

IOO 

o.  2241 

—  0.0003 

O.2244 

0.0048 

0-54 

4 

2OO 

0.2248 

+O.OOO4 

O.2244 

0.0048 

2.44 

4 

200 

0.2245 

+0.0001 

The  Determination  of  Mercury  by  Titration  with  Sodium  Thio- 

sulphate.  •  *. 

The  method  proposed  by  Scherer*  for  the  determination  of 
mercury  by  titration  of  mercury  salts  with  sodium  thiosulphate 
has  been  studied  by  Norton. f  It  is  shown  in  the  estimation  of 
mercury  taken  in  the  form  of  mercuric  chloride  the  reaction  pro- 
ceeds definitely  according  to  the  equation 

3  HgCl2+2  Na2S203+3  H2O  =  2  HgS.HgCl,+2  Na2SO4+4  HC1. 

*  Lehrbuch  der  Chemie,  i,  513. 

t  John  T.  Norton,  Am.  Jour.  Sci.,  (4],  x,  48. 


MERCURY 


197 


This  method  yields  accurate  results  if  carried  out  under  certain 
fixed  conditions.  These  conditions,  which  must  be  closely  ad- 
hered to,  are  as  follows:  The  solution  containing  the  mercury 
in  the  form  of  mercuric  chloride  is  diluted  to  100  cm.3  and  heated 
to  a  temperature  of  60°  C.  The  sodium  thiosulphate  in  ^Oth 
normal  solution  is  run  in  from  a  burette  until  the  white  pre- 
cipitate, 2  HgS.HgCl2,  first  formed  begins  to  take  on  a  brownish 
tinge  due  to  incipient  formation  of  the  sulphide.  The  solution  is 
then  diluted  with  cold  water,  some  asbestos  fiber  added  to  coagu- 
late the  precipitate,  and  the  whole  is  quickly  thrown  on  the  filter. 
After  careful  washing,  potassium  iodide  is  added  to  the  filtrate 
and  the  excess  thiosulphate  is  titrated  with  iodine  in  presence 
of  starch.  The  duration  of  the  process  need  not  exceed  fifteen 
minutes.  It  is  worthy  of  note  that  there  is  no  necessity  of  using 
any  hydrochloric  acid  in  addition  to  that  formed  in  the  reaction. 
Results  obtained  by  this  procedure  are  given  in  the  table. 

Titration  of  Mercuric  Chloride. 


HgCl2  taken 
as  Hg. 

grm. 

Volume  at 
beginning. 

cm.3 

Temper- 
ature. 

C. 

Na2S203  in 
excess. 

cm.a 

HgCl2  found 
as  Hg. 

grill* 

Error, 
grm. 

0.0759 

TOO 

60° 

3-06 

0.0766 

+0.0007 

0.0384 

IOO 

60° 

2.8l 

0.0387 

+0  .  0003 

o.  1498 

IOO 

60° 

I  .  I 

0.1500 

+0.0008 

0.1503 

IOO 

60° 

1.63 

0.1506 

+o  .  0003 

0.1479 

IOO 

60° 

2.41 

o.  1480 

+O.OOOI 

o.  1489 

IOO 

60° 

2.  12 

0.1503 

+0.0014 

o.  2244 

IOO 

60° 

2.63 

0.2259 

+0.0015 

o.  1490 

IOO 

60° 

2-33 

0.1484 

—0.0006 

0.0758 

IOO 

60° 

2. 

0.0762 

+o  .  0004 

0.0383 

IOO 

60° 

2.53 

0.0379 

—0.0004 

In  applying  Scherers  process  to  mercurous  nitrate  and  mer- 
curic nitrate,  Norton  was  unable  to  discover  conditions  of  definite 
action. 

The  Estimation  of  Mercury  by  Precipitation  as  Mercurous  Oxalate 
and  Titration  of  the  Excess  of  Precipitant  with  Permanganate. 

Peters*  has  shown  that  mercury  taken  as  mercurous  nitrate 
may  be  estimated  by  precipitating  it  as  the  oxalate  and  deter- 
mining by  titration  with  potassium  permanganate  the  excess  of 

*  C.  A.  Peters,  Am.  Jour.  Sci.,  [4],  ix,  401. 


198 


METHODS  IN  CHEMICAL  ANALYSIS 


ammonium  oxalate  used  as  the  precipitant.*  The  operation  is 
successful  in  presence  of  2  per  cent  to  5  per  cent  of  dilute  nitric 
acid,  sp.  gr.  1.15,  in  100  cm.3  of  the  mixture;  and  mercuric  salt 
to  the  extent  of  10  per  cent  of  the  amount  of  the  mercurous  salt 
may  be  present  provided  the  nitric  acid  amounts  to  2  per  cent, 
and  the  excess  of  ammonium  oxalate  used  to  effect  precipitation 
is  not  too  great.  The  precipitate  settles  well  when  properly 
stirred,  and  the  excess  of  the  precipitant  need  not  exceed  the 
safe  limit  of  I  cm.3  or  2  cm.3.  Following  are  results  obtained  by 
the  permanganate  titration  of  the  filtrate  from  the  precipitated 
mercurous  oxalate. 

Precipitation  by  Ammonium  Oxalate  and  Titration  of  the  Excess. 


Hg  taken  as 
Hg2(N03)2. 

grin. 

Hg  present 
as  Hg(NO,),. 

grin. 

Excess  of 
w/io  ammo- 
nium oxalate. 

cm.1 

HN03 

(sp.  gr. 
I.IS). 

cm.1 

Volume  at 
precipita- 
tion. 

cm.8 

Hg  found. 

gnu  . 

Error. 

O.I2I7 

0.0067 

0.86 

2 

IOO 

0.1232 

—0.0015 

0.1217 

0.0067 

0.92 

2 

IOO 

O.  I22O 

+0.0003 

o.  1217 

0.0067 

0.97 

2 

IOO 

0.1218 

-j-o.oooi 

o.  1217 

0.0067 

0.90 

2 

IOO 

o.  1224 

+0.0007 

O.I2I7 

0.0067 

0.91 

2 

IOO 

0.1213 

—0.0004 

O.I22I 

0.0067 

3-92 

2 

IOO 

0.1237 

+0.0016 

O.I2I7 

0.0067 

3-93 

2 

IOO 

0.1235 

+0.0018 

O.I242 

0.0067 

8-75 

2 

IOO 

o.  1263 

+O.OO2I 

o.  1217 

0.0134 

0.87 

2 

IOO 

0.1230 

+0.0013 

O.I2I7 

0.0134 

0.86 

2 

IOO 

0.1232 

+0.0015 

O.I2I7 

0.0134 

3-93 

4 

IOO 

o.  1218 

+O.OOOI 

O.I2I7 

0.0134 

3-90 

4 

IOO 

O.I2II 

—0.0006 

0.2244  . 

0.0067 

1.88 

4 

115 

0.2244 

o  .  oooo 

O.2244 

0.0067 

1.91 

4 

130 

0.2240 

—  0.0004 

0.2244 

0.0141 

2.98 

4 

IOO 

0.2230 

—  0.0014 

0.2244 

0.0141 

2.94 

4 

IOO 

0.2241 

—0.0003 

The  Titration  of  Mercurous  Salts  with  Potassium  Permanganate. 

If  a  solution  of  a  mercurous  salt,  such  as  mercurous  sulphate, 
in  dilute  sulphuric  acid,  is  titrated  with  potassium  permanganate 
in  the  usual  manner,  the  bleaching  of  the  color  is  rapid  at  first, 
but  long  before  the  oxidation  is  complete  the  solution  assumes 
a  golden-yellow  color  and  on  standing  the  brown  oxides  of  man- 
ganese are  precipitated.  For  this  reason  no  definite  end  reac- 
tion is  obtainable.  This  difficulty  may,  however,  be  avoided,  as 
Randall  t  has  shown,  if  the  permanganate  solution  is  added  in 

*  See  page  195. 

f  D.  L.  Randall,  Am.  Jour.  Sci.,  [4],  xxiii,  137. 


MERCURY 


199 


excess,  the  color  then  bleached  with  a  standard  ferrous  sulphate 
solution,  and  the  end-point  finally  reached  by  a  few  drops  of 
permanganate.  Under  these  conditions  the  end  reaction  is  per- 
fectly sharp  and  the  oxidation  of  the  mercurous  salt  complete. 

According  to  Randall's  procedure,  the  mercurous  sulphate  or 
nitrate  in  solution  in  water  containing  a  suitable  excess  of  sul- 
phuric acid,  or  nitric  acid  free  from  oxides  of  nitrogen,  is  diluted 
and  treated  with  n/2O  potassium  permanganate  until  the  mix- 
ture, colored  brown  by  the  oxides  of  manganese,  takes  on  a 
distinctly  red  tint.  Ferrous  sulphate  in  n/io  solution  is  added 
in  amount  sufficient  to  clear  the  solution,  and  the  titration  is 
immediately  completed  with  permanganate.  The  n/2O  solu- 
tions are  preferred  to  the  usual  n/io  solutions  on  account  of  the 
high  equivalent  of  the  mercurous  salt.  Even  with  n/2O  solu- 
tion, o.i  cm.3  of  the  permanganate  is  equivalent  to  about  o.ooio 
grm.  of  mercury.  For  the  same  reason  the  titrations  are  made 
with  all  possible  care  and  accuracy.  Following  are  experimental 
results  obtained  with  a  mercurous  sulphate  solution  prepared  by 
shaking  up  an  excess  of  the  salt  in  water  acidified  with  sulphuric 
acid,  allowing  the  mixture  to  stand  twenty-four  hours,  and 
filtering  through  asbestos.  The  solution  was  standardized  by 
weighing  the  mercurous  chloride  precipitated  by  sodium  chlo- 
ride and  dried  in  a  vacuum. 


Titration  of  Mercurous  Sulphate. 


Hg2S04 
sol. 

cm.* 

H2S04 
[i  :  i]. 

cm.*  ' 

KMn04 
approx. 
n/2o. 

cm.8 

FeSO4  = 
cm.8 

=  KMnO4 
cm.8 

KMnO4 
final. 

cm.8 

Hg 

found. 

grm. 

Hg. 
(theory). 

grm. 

Error, 
grm. 

Volume  150  cm.3. 


IOO 

5 

14.70 

10 

10.90 

3.80 

0.0346 

0-0354 

—0.0008 

IOO 

5 

14.72 

10 

10.90 

3.82 

0.0347 

0.0354 

—0.0007 

IOO 

5 

14.70 

10 

10.90 

3.80 

0.0346 

0.0354 

—0.0008 

IOO 

5 

14.71 

10 

10.90 

3.81 

0.0346 

0.0354 

—0.0008 

Hg2S04 
sol. 

Volume  500  cm.3. 

505.8 

5 

29.20 

IO 

10.90 

18.30 

0.1668 

0.1666 

+O.OOO2 

500.2 

5 

32.16 

13 

14.17 

17.99 

0.1639 

0.1648 

—  0.0009 

510-1 
499-2 

5 
5 

29.29 
28.95 

IO 
IO 

10.90 
10.90 

18.39 
18.05 

0.1676 

0.1645 

0.1680 
0.1644 

—  0.0004 
+0.0001 

200 


METHODS  IN  CHEMICAL  ANALYSIS 


For  practical  purposes  the  application  of  the  method  to  mer- 
curous  nitrate  is  of  much  greater  importance.  Results  obtained 
with  mercurous  nitrate  in  solution  —  prepared  by  dissolving  the 
crystallized  salt  in  water  containing  enough  pure  nitric  acid  to 
prevent  the  formation  of  basic  salts,  and,  as  an  additional  pre- 
caution, passing  a  current  of  hydrogen,  washed  by  alkaline  per- 
manganate and  alkaline  pyrogallol,  for  twelve  hours  through  the 
solution  to  remove  nitrous  acid  —  are  given  below. 

Titration  of  Mercurous  Nitrate. 
Volume  200  cm.3. 


TTa 

(NC&2 

HNO3. 

Approx. 

H/2O 

KMn04. 

FeSO4.  = 

KMnO4 

KMnO4 
final. 

Hg 
found. 

Hg 
theory. 

Error. 

cm.3 

cm.3 

cm.3 

cm.8 

cm.3 

cm.3 

grm. 

grm. 

grm. 

25 

5 

49-99 

9.^ 

10.64 

39-35 

0.3586 

0-3594 

—  O.OOoS 

25 

5 

49-68 

9-50 

10.36 

39-32 

0.3583 

0-3594 

—  O.OOII 

25 

5 

49-73 

9-50 

10.36 

39-37 

0.3588 

0-3594 

—  O.O006 

25 

5 

49.89 

9.67 

10.54 

39-35 

0.3586 

0-3594 

—  O.OOOS 

25 

5 

49.70 

9-50 

10.36 

39-34 

0.3585 

0-3594 

—  O.OOO9 

H2S04 

[i  :  i]. 

25 

5 

50.29 

IO.OO 

10.90 

39-39 

0.3589 

0-3594 

—  0.0005 

CHAPTER  VI. 
BORON;  ALUMINIUM;  LANTHANUM;  THALLIUM. 

BORON. 

The  Gravimetric  Determination  of  Boric  Acid. 

The  use  of  ^HE  Process  °f  determining  boric  acid  by  distil- 

Caicium Oxide  lation  with  methyl  alcohol,  evaporation  of  the  dis- 
as  a  Retainer.  tjnate  on  calcium  oxide,  and  ignition  of  the  residue,* 
has  been  reexamined  by  Gooch  and  Jones  f  in  the  light  of  the 
experience  of  many  investigators.t  It  is  shown  that  difficulties 
arising  when  nitric  acid  is  present  in  the  retort  may  be  obviated 
by  limiting  the  amount  of  that  acid  by  the  use  of  phenolphthalein 
as  an  indicator  at  the  outset  of  the  distillation.  The  addition  of 
a  drop  of  the  acid  and  another  of  the  indicator  should  be  re- 
peated once  or  twice  during  the  distillation  to  insure  the  replace- 
ment of  the  acid  volatilized  from  the  salt  slightly  decomposed 
in  the  process.  The  effect  of  much  nitric  acid  is  bad,  not  only 
because  it  neutralizes  the  calcium  oxide  when  it  passes  to  the 
distillate,  but  because  when  it  is  used  the  dried  mixture  of 
calcium  hydroxide  and  borate  containing  nitrate,  nitrite  and 
organic  matter  is  likely  to  puff  explosively  if  ignition  is  begun  as 
soon  as  the  residue  is  dry.  If  the  residue  is  heated  gradually 
and  as  strongly  as  possible  over  a  radiator  before  the  flame  is 
actually  applied  to  the  crucible,  no  such  action  takes  place. 

That  good  results  may  be  obtained  with  small  amounts  of 
calcium  oxide,  provided  care  as  to  the  use  of  nitric  acid  and  the 
conditions  of  ignition  be  taken,  is  shown  by  the  figures  of  the 
original  description  and  by  the  following  results  of  experiments 
in  which  phenolphthalein  was  employed  as  an  indicator  and 
the  residue  heated  strongly  over  the  radiator  before  actual 
ignition. 

*  Gooch,  Am.  Chem.  Jour.,  ix,  23. 

t  F.  A.  Gooch  and  Louis  Cleveland  Jones,  Am.  Jour.  Sci.,  [4],  vii,  34. 
t  Penfield,  Am.  Jour.  Sci.,  [3],  xxxiv,  222;  Kraut,  Zeit.  anal,  Chem.,  xxxvi, 
3;  Moissan,  Compt.  rend.,  cxvi,  1084. 

201 


202 


METHODS  IN  CHEMICAL  ANALYSIS 


Distillation  with  Nitric  Acid. 


CaO  taken. 

B2O3  taken. 

B2O3  found. 

Error. 

grin. 

grm. 

grm. 

grm. 

2.3405 
i  .  7620 
2.1757 
2.5656 

0.1788 
0.1790 
».            0.1824 
0.1788 

0.1792 
0.1785 
0.1840 
0.1786 

+9.0004 
—  0.0005 
+0.0016 

—  O.OO02 

No  difficulty  exists  in  respect  to  explosive  effects  when  acetic 
acid  is  used  in  place  of  nitric  acid,  though  even  in  this  case  it  is 
safer  to  use  the  radiator  in  the  first  stages  of  heating,  thus  avoid- 
ing the  danger  of  mechanical  loss  by  too  rapid  ignition. 

Following  are  determinations  made  by  this  method  with  the 
use  of  acetic  acid. 

Distillation  with  Acetic  Acid. 


CaO  taken. 

B2O3  taken. 

B2Oa  found. 

Error. 

grm. 

grm. 

grm. 

grm. 

0.9977 
I.O22O 

o  .  2065 

0.2067 

O  .  2062 
0.2070 

—  0.0003 
+0.0003 

I-37I7 
I.I3IO 

0.2077 
0.1791 

0.2075 
0.1795 

—  O.OOO2 
+O.OOO4 

The  results  of  the  preceding  table,  as  well  as  those  of  the  inves- 
tigators mentioned,  are  a  sufficient  answer  to  the  criticism*  that 
acetic  acid  and  nitric  acid  do  not  liberate  boric  acid  in  the  dis- 
tillation process  so  that  good  results  may  be  obtained.  In  fact, 
it  has  been  shown  f  that  the  prolonged  action  of  carbonic  acid  is 
adequate  to  bring  about  complete  volatility  of  boric  acid  with 
methyl  alcohol.  The  number  of  distillations  required  depends, 
of  course,  upon  the  amount  of  boric  acid  to  be  volatilized.  To 
remove  0.2  grm.  of  the  anhydride  completely  to  the  distillate, 
it  was  shown  in  the  original  description  of  the  method  J  that  six 
io-cm.3  charges  of  methyl  alcohol  suffice. 

It  is  also  shown  that  the  difficulty  §  in  bringing  calcium  oxide 
to  a  constant  weight  before  and  after  absorption  of  boric  acid 
has  been  magnified  unduly.  Thus  the  following  table  shows  the 

*  Reischle,  Zeit.  anal.  Chem.,  xxvi,  512. 

f  Jones,  Am.  Jour.  Sci.,  [4],  v,  442. 

J  Loc.  cit. 

§  Thaddeef,  Zeit.  anal.  Chem.,  xxxvi,  .568. 


BORON 


203 


series  of  weights  taken  in  several  experiments  in  bringing  calcium 
oxide  to  a  constant  weight  in  a  5O-cm.3  platinum  crucible  ignited 
over  a  blast  lamp,  as  well  as  the  weight  taken  after  adding  a 
known  amount  of  standard  boric  acid  solution  to  the  slaked 
oxide,  evaporating,  and  igniting.  The  results  recorded  are  those 
of  experiments  made  on  days  not  moist  beyond  the  average,  and 
with  the  greatest  care  to  approach  the  limit  of  accuracy  with 
which  calcium  oxide  and  the  boric  acid  held  thereby  can  be 
weighed  under  ordinarily  favorable  conditions.  The  first  weight 
of  calcium  oxide  recorded  under  each  experiment  was  taken  after 
a  strong  ignition  over  the  blast  lamp  for  about  one-half  hour. 
The  succeeding  weights  were  taken  after  similar  ignition  of  five 
minutes.  In  all  cases  the  crucible  was  left  to  stand  a  definite 
period  in  a  sulphuric  acid  dessicator,  and,  after  the  approximate 
value  had  once  been  obtained,  the  weights  of  the  preceding  weigh- 
ing were  replaced  on  the  balance  before  the  crucible  was  taken 
from  the  dessicator.  The  average  of  the  weights  bracketed  is 
the  weight  taken  as  constant  for  the  calculations. 

Ignition  of  Calcium  Oxide  and  Calcium  Bofate. 


CaO  taken, 
grm. 

B2O3  taken, 
grm. 

CaO+B2O3 
taken. 

grm. 

CaO+B203 
found. 

grm. 

Error, 
grm. 

(0.9505 

(i  Ho.  9493  [o.9493 
(0.9493  \ 

0.2095 

I  .  1588 

i'::S?}'-'*» 

+0.0003 

(4;.lii'-1315 

0.2150 

L3465 

1-3499 
1-3474) 
1.3475  >  1.3475 
1.3476  ) 

+O.OOIO 

(0.8028 

0.9205 

(3)  (0:8024  1°-8°24 

0.1184 

0.9208 

0^206  I0"9206 

+0.0002        1 

C  2  .  6980 

(4)J  2-6975 

i    2    6o73   I         f 

U.'  6973  12-6973 

0.2073 

2.9046 

2  .  9043 

+O.OOO2 

Obviously  calcium  oxide  may  be  weighed  with  accuracy,  with 
or  without  boric  acid ;  but  the  fact  remains  that  a  less  hygroscopic 
absorbent  —  one  requiring  less  care  in  the  handling  —  is  to  be 
desired. 


204  METHODS  IN  CHEMICAL  ANALYSIS 

The  Use  of  Gooch  and  Jones  find  that  sodium  tungstate,  fused 


1™8"  a  sn'gnt  excess  of  tungstic  acid  over  that  con- 


Retainer.  tained  in  the  normal  tungstate  (to  insure  its  freedom 

from  carbonate),  may  be  used  with  good  results  as  an  absorbent 
for  boric  acid.  This  substance  is  definite  in  weight,  not  hydro- 
scopic,  soluble  in  water,  and  recoverable  in  its  original  weight 
after  evaporation  and  ignition. 

According  to  the  procedure  advocated,  use  is  made  of  the 
apparatus  originally  proposed,  so  modified  that  the  Erlenmeyer 
flask  used  as  a  receiver  is  fitted  tightly  to  the  condenser  and 
trapped  with  water  bulbs.*  The  retort  is  made  very  easily  from 
a  i5O-cm.3  pipette  and  has  the  special  advantage  that  particles 
of  the  residue  spattering  during  distillation  are  easily  washed 
from  the  walls  of  the  vessel  by  a  slight  rotary  motion  of  the 
retort.  Special  care  should  be  taken  to  give  the  tungstate  ample 
time  for  contact  with  the  distillate  before  exposing  the  latter  to 
atmospheric  evaporation.  The  distillate  is  received,  therefore, 
in  a  dilute  solution  of  sodium  tungstate  which  is  placed  in  the 
receiver  cooled  by  ice  and  trapped  with  water.  The  mixture  is 
well  stirred,  allowed  to  stand  one-half  hour  after  the  distilla- 
tion is  complete,  evaporated  to  small  volume  in  a  large  dish,  and 
transferred  to  the  crucible  in  which  the  tungstate  was  originally 
weighed.  After  thorough  drying  the  residue  is  ignited  to  fusion 
and  weighed.  When  acetic  acid  is  employed  in  the  retort,  care 
must  be  taken  in  the  ignition  to  expose  the  fused  mass  freely  to 
the  air  (by  causing  it  to  flow  upon  the  sides  of  the  crucible)  until 
the  color  of  the  cooled  tungstate  is  white,  in  order  that  the  re- 
ducing effect  of  the  acetate  may  be  eliminated.  In  the  experi- 
ments recorded  in  the  following  table  the  tungstate  was  used  in 
the  receiver  to  retain  the  boric  acid,  distilled  as  usual  with  methyl 
alcohol  from  the  borates  treated  with  acetic  acid,  nitric  acid  or 
sulphuric  acid,  in  amounts  regulated  by  the  use  of  phenolphtha- 
lein  as  an  indicator.  Excessive  use  of  acid  is  disadvantageous, 
and  this  is  especially  true  in  the  case  of  sulphuric  acid;  for,  if 
this  acid  is  carried  over  with  the  methyl  alcohol,  as  it  is  at  100° 
if  present  in  appreciable  excess,  a  part  of  it,  at  least,  is  held  per- 
manently by  the  tungstate  to  increase  the  apparent  weight  of 
the  boric  acid  estimated.  The  number  of  distillations  necessary 
depends  upon  the  amount  of  boric  acid  to  be  volatilized.  Six 
*  See  Fig.  2,  page  3. 


BORON 


205 


charges  of  10  cm.3  each  of  methyl  alcohol  are  enough  to  transfer 
0.2  grm.  of  boric  anhydride  to  the  distillate. 

Distillation  of  Six  10  cm.3  Portions  of  Methyl  Alcohol. 


Na2WO4+WO3  taken. 

B2O3  taken. 

B2O3  found. 

Error. 

gnu. 

grm. 

grm. 

grm. 

With  nitric  acid! 

8.5516 

4-9639 
8.0033 

0.1582 
0.1329 
0.1267 

0.1572 
0.1323 
o.  1256 

—  O.OOIO 
—  O.OOO6 
—  O.OOII 

With  acetic  acid. 


4.9658 

0.1434 

o.  1418 

—  0.0016 

6.0289 

0.1431 

0.1433 

4-O.OOO2 

4.6797 

0.1589 

0.1587 

—  O.OOO2 

4.0013 

0.1433 

0.1422 

—  O.OOII 

With  sulphuric  acid. 


6.3439 

0.1582 

0.1579 

-0.0003 

8.8227 

0.1582 

0.1577 

-0.0005 

10.  1516 

0.1265 

o.  1264 

—  O.OOOI 

6.5738 

0.1392 

0.1390 

—  O.OOO2 

The  Acidimetric  Estimation  of  Boric  Acid. 
When  boric  acid  and  mannite  are  mixed  in  solution,  a  peculiar 
compound  of  strongly  acid  properties  is  the  result.  This  com- 
pound decomposes  carbonates,  and  its  acid  taste  is  comparable 
to  that  of  citric  acid,  much  stronger  than  that  of  boric  acid  alone. 
Magnanini  *  has  found  that  the  product  of  such  a  mixture  of  boric 
acid  and  mannite  solutions  shows  greater  electrical  conductivity 
and  a  lower  freezing  point  than  a  similar  molecular  solution  of 
either  substance  alone.  Other  polyatomic  alcohols  (but  all  to  a 
less  degree  than  mannite)  and  some  organic  acids  show  this 
peculiar  property  of  combining  chemically  with  boric  acid  to 
increase  its  acid  qualities. f  Of  the  reaction  between  boric  acid 
and  glycerin,  Thomson,  t  Barthe,§  and  J6rgensen||  have  taken 

*  Gaz.  Chim.,  xx,  428-440;  xxi,  134-145. 

f  Klein,  Compt.  rend.  Ixxxvi,  826;  xcix,  144.     Lambert,  ibid.,  cviii,  1016— 
1017. 

t  J.  S.  C.  I.,  xv,  432. 

§  J.  Pharm.  Chim.,  xxix,  163. 

II  Zeit.  angew.  Chem.  (1897),  5. 


206  METHODS  IN  CHEMICAL  ANALYSIS 

advantage  to  develop  methods  for  the  volumetric  estimation  of 
boric  acid,  glycerin  being  used  to  form  a  combination  with  boric 
acid  sufficiently  acidic  to  give  an  acid  reaction  when  used  with 
a  sensitive  indicator  and  make  possible  its  titration  with  an  alkali 
solution;  but  Honig  and  Spitz*  show  that  (in  the  method  of 
Jorgensen)  a  very  large  amount  of  glycerin  must  be  used  to  pre- 
vent the  appearance  of  the  indication  of  alkalinity  with  phenol- 
phthalein  before  all  the  boric  acid  is  neutralized  according  to  the 
equation  2  NaOH  +  B2O3  =  2  NaOBO  +  H2O.  Vadam  f  uses 
mannite,  which,  as  he  finds,  gives  sharper  indications  with 
litmus.  The  solution  must  be  boiled  to  decompose  bicarbonates 
while  the  volatilization  of  boric  acid  is  prevented  by  the  use  of  a 
return  condenser;  and  silica  must  be  removed  by  the  process  of 
Berzelius,  and  the  solution  neutralized  to  stronger  acids,  before 
a  titration  of  the  boric  acid  can  be  made. 

Many  indicators  said  to  be  insensible  to  free  boric  acid  have 
been  used  to  indicate  the  neutralization  of  the  stronger  acids. 
Honig  and  Spitz, J  and  Thomson, §  use  methyl  orange;  Morse  and 
Burton||  use  tropaeolin  oo;  while  Vadam  IT  makes  use  of  litmus. 
Neutralization  Finding  all  these  indicators  to  be  more  or  less 
of  stronger  affected  by  boric  acid,  Jones**  has  had  recourse  to  the 
well-known  reaction  according  to  which  a  stronger 
acid  liberates  regularly  iodine  from  a  mixture  of  iodide  and  iodate, 
which  is  the  solution  of  this  difficulty.  If  both  the  iodide  and 
iodate  are  in  excess  of  the  acid,  the  entire  amount  of  free  acid  will 
be  neutralized  and  the  corresponding  amount  of  iodine  liberated 
according  to  the  following  equation : 

5  KI  +  KIO3  +  6  HC1  =  6  KC1  +  3  H2O  +  3  I2. 

This  liberated  iodine  may  be  removed  by  sodium  thiosulphate 
and  a  solution  obtained  which  is  absolutely  neutral,  containing 
potassium  iodide,  iodate  and  tetrathionate.  The  statements 
made  by  P.  Georgevicft  and  Furry, jj  that  boric  acid  present  in 

*  Zeit.  angew.  Chem.  (1896),  549. 

t  J.  Pharm.  Chim.  [6],  viii,  109-111. 

J  Zeit.  anorg.  Chem.,  xviii,  549. 

§  J.  S.  C.  I.,  xv,  432. 

||  Am.  Chem.  Jour.,  x,  154. 
1  J.  Pharm.  Chim.  [6],  viii,  109-111. 
**  L.  C.  Jones,  Am.  Jour.  Sci.,  [4],  vii,  147. 
tt  J«  prakt.  Chem.,  xxxviii,  118. 
tt  Am.  Chem.  Jour.,  vi,  341. 


BORON  207 

moderate  amount  in  solution  has  not  the  slightest  action  on  a 
mixture  of  iodide  and  iodate,  have  been  found  to  apply  to  solu- 
tions containing  not  more  than  o.i  grm.  of  B2C>3  to  25  cm.3. 
Therefore,  when  this  acid  is  liberated  by  an  excess  of  a  stronger 
acid  and  the  iodine  set  free  destroyed  by  thiosulphate,  it  remains 
free  in  solution  to  be  titrated  in  any  convenient  manner.  Follow- 
ing along  the  lines  suggested  by  the  above  reactions,  Jones  *  has 
developed  a  volumetric  process  for  the  estimation  of  boric  acid. 

The  solution  in  which  boric  acid  is  present  to  an  amount  not 
exceeding  o.  I  grm.  in  25  cm.3  is  made  slightly  acid  to  litmus  by 
hydrochloric  acid  and  treated  with  5  cm.3  of  a  solution  (10  per 
cent)  of  barium  chloride.  An  amount  of  iodate  and  iodide  of 
potassium  sufficient  to  liberate  iodine  at  least  equivalent  to  the 
excess  of  hydrochloric  acid  in  the  acidified  solution  is  mixed  with 
starch  in  a  separate  beaker,  the  iodine  which  is  usually  thrown 
out  by  this  mixture  being  just  bleached  by  a  dilute  solution  of 
thiosulphate. 

To  the  now  neutral  solution  of  iodide  and  iodate  a  single  drop 
of  the  solution  to  be  analyzed  is  transferred  by  a  glass  rod.  If 
a  blue  coloration  is  developed,  the  solution  is  known  to  be  acidic 
with  hydrochloric  acid,  and  all  the  boric  acid  is  in  free  condition. 
The  amount  of  iodide  and  iodate  necessary  depends  upon  the 
acidity  of  the  solution  containing  boric  acid.  Usually  a  mixture 
of  10  cm.3  of  a  25  per  cent  solution  of  iodide  with  the  same 
amount  of  a  saturated  solution  of  iodate  is  sufficient.  Any  large 
excess  of  hydrochloric  acid  should  be  neutralized  by  sodium 
hydroxide  before  the  iodide  and  iodate  mixture  is  added.  After 
the  addition  of  the  iodide  and  iodate  solution  and  starch  to  the 
boric  acid  solution,  the  liberated  iodine  is  carefully  bleached  by 
thiosulphate.  Excess  of  thiosulphate  in  reasonable  amount  does 
not  seem  to  be  detrimental,  but  in  practice  the  starch  iodide  color 
is  clearly  bleached,  and  then  no  more  is  added. 

Soluble  carbonates  prevent  a  definite  indication  of  the  neutral 
point  by  thiosulphate  and  starch  iodide,  therefore  the  barium 
chloride  was  added  to  transform  them  to  insoluble  barium  car- 
bonate in  the  action  of  the  iodide-iodate  mixture.  The  mixture 
of  iodide  and  iodate  is  not  added  to  the  solution  to  be  analyzed 
until  after  it  has  been  made  distinctly  acidic,  for  the  reason 
that,  when  the  neutral  point  is  approached  in  the  addition  of 
*  Am.  Jour.  Sci.,  [4],  viii,  129. 


208  METHODS  IN   CHEMICAL  ANALYSIS 

hydrochloric  acid,  the  starch  iodide  thrown  out  locally  by  the  acid 
is  not  bleached  again  by  the  small  amount  of  sodium  borate 
remaining  undecomposed  and  thus  obscures  the  neutral  point, 
strengthening  The  solution,  after  the  bleaching  of  iodine  by 
of  Boric  Acid  thiosulphate,  is  colorless  and  contains  only  starch, 
neutral  chloride,  potassium  tetrathionate,  iodide  and 
iodate,  and  all  the  boric  acid  present  in  uncombined  condition. 
The  carbonate  lies  out  of  the  sphere  of  action  in  insoluble  form 
as  barium  carbonate.  A  few  drops  of  the  indicator,  phenol- 
phthalein,  are  now  added,  and  n/$  sodium  hydroxide  is  run  in 
until  a  strong  red  coloration  is  produced.  A  pinch  of  mannite, 
I  grm.  or  2  grm.,  is  then  added,  which  bleaches  the  phenol- 
phthalein  coloration,  and  the  alkali  solution  again  run  in  to  a 
faint  indication,  which,  if  permanent  on  the  addition  of  more 
mannite,  may  be  taken  as  the  reading  point. 

The  combination  of  boric  acid  and  mannite  liberates  immedi- 
ately in  the  presence  of  iodide  and  iodate  about  half  the  amount 
of  iodine  required  on  the  theory  that  B2O3  acts  with  the  neu- 
tralizing power  of  metaboric  acid,  HOBO.  If  no  mannite  is 
present  phenolphthalein  gives  an  alkaline  indication  when  only 
about  one-half  the  amount  of  alkali  theoretically  necessary  to 
form  the  metaborate,  NaOBO,  has  been  added.  Obviously, 
then,  the  starch  iodide  coloration  will  not  appear  on  the  addition 
of  mannite,  if  the  free  boric  acid  has  been  neutralized  to  phenol- 
phthalein by  alkali,  and  the  remainder  of  the  alkali  is  added 
to  complete  neutralization  immediately  after  the  addition  of 
mannite.  The  end  reaction  with  phenolphthalein  is  sharp  and 
the  small  amount  of  carbonate  present  in  the  standard  solution 
of  alkali  is  precipitated  by  the  barium  chloride  already  in  the 
solution.  The  calculation  must,  therefore,  be  based  on  the 
amount  of  free  alkali  hydroxide  used,  according  to  the  following 
representation : 

B2O3  +  2  NaOH  =  2  NaOBO  +  H2O. 

The  best  results  and  the  most  definite  indications  are  obtained 
in  cold  solution  of  a  volume  not  greater  than  50  cm.3.  This  fact 
accords  with  the  observations  of  Magnanini*  that  the  relative 
electrical  conductivity  of  the  boromannite  solution  is  decreased 
by  dilution  and  elevation  of  the  temperature.  When  silicates  are 
*  Gaz.  Chim.,  xx,  428,  and  xxi,  134. 


BORON 


209 


present  in  solution,  the  silicon  dioxide  is  liberated  by  the  excess 
of  hydrochloric  acid,  and  this  oxide,  whether  in  hydrous  or  an- 
hydrous condition,  neither  affects  the  indication  with  iodine  nor 
phenolphthalein,  nor  does  it  form  with  mannite  a  compound  of 
acidic  properties.  The  presence  of  fluorides  is  not  detrimental. 
Ammonium  salts  do  interfere  with  the  indication  given  by 
phenolphthalein,  but  they  may  be  removed  by  boiling  with 
potassium  hydroxide  in  excess,  or  an  indicator  which  is  not 
affected  by  them  may  be  used. 

The  following  table  contains  the  results  of  a  series  of  analyses 
in  which  the  boric  acid  was  first  drawn  into  an  excess  of  sodium 
hydroxide,  then  estimated  according  to  the  method  described. 

Titration  of  Boromannite  Solution  with  Standard  Alkali. 


B2O3  sol.  taken, 
cm.' 

NaOH  sol. 
required. 

cm.3 

B2OS  taken, 
grm. 

B2O3  found, 
grm. 

Errors  on  B2O3. 
grm. 

21-95 

21.  02 

0.1571 

0.1577 

+o  .  0006 

20.68 

19.65 

0.1479 

0.1474 

—  0.0005 

20.73 

19.63 

0.1483 

0.1473 

—  O.OOIO 

23-05 

23.71 

0.1776 

0.1777 

+0.0001 

23.10 

23.80 

0.1780 

0.1783 

+0.0003 

22.76 

23-35 

0.1754 

0.1750 

—0.0004 

24.08 

24.78 

0-1855 

0.1857 

+O.OOO2 

22.00 

22.50 

0.1695 

0.1686 

—  O.OOOQ 

20.78 

21.28 

0.1601 

0.1595 

—  O.OOO6 

Practical  tests  of  the  method  upon  crude  calcium  borate  and 
colemanite  are  given  below.  The  finely  ground  minerals  were 
dissolved  in  hydrochloric  acid  and  the  analyses  proceeded  with 
as  above  described. 

A  determination  of  boric  acid  by  this  process  can  be  completed 
in  five  minutes  and  the  results  are  obviously  accurate  within 
the  limits  of  ordinary  analysis. 

Analysis  of  Crude  Borate  of  Lime. 


Ca  borate  taken. 

B2O3  found. 

B203. 

grm. 

grm. 

Per  cent. 

0.4016 
0.4044 
0.4000 

0.2289 
0.2302 
0.2285 

56.99 
56.92 

57-11 

210 


METHODS  IN   CHEMICAL  ANALYSIS 


Analysis  of  Colemanite. 


Mineral  taken, 
grm. 

B2O,  found, 
grm. 

B203. 
Per  cent. 

Average. 
Per  cent. 

0.4034 

O  .  2064 

51.15 

] 

o  .  4070 

o  .  2069 

50.80 

0.6004 
0.6006 

0.3054 
0.3056 

50.86 
50.89 

|-            50.99 

0-5059 

0.2592 

5I-24 

0.5092 

0.2592 

50.89 

J 

The  usually  interfering  substances,  fluorine,  silica  and  car- 
bon dioxide,  have  no  detrimental  influence  on  the  results  of  this 
process. 

The  lodometric  Determination  of  Boric  Acid. 

In  studying  the  strongly  acidic  compound  formed  when  boric 
acid  and  mannite  are  associated  in  solution,  Jones*  finds  that 
the  acid  developed  is,  under  certain  definite  conditions,  suffi- 
ciently strong  to  liberate,  quantitatively,  from  a  mixture  of 
potassium  iodide  and  iodate,  the  amount  of  iodine  required  on 
the  supposition  that  each  molecule  of  metaboric  acid  (HOBO) 
acts  in  a  manner  similar  to  a  univalent  mineral  acid  under  the 
same  conditions. 

5KI  +KIO3  +  6HOBO  =  3  I2  +  6  KOBO  +  3  H2O. 

Obviously,  this  reaction  depends  upon  the  behavior  of  the 
acidic  boromannite  compound  as  an  acid  stronger  than  acetic, 
tartaric  or  citric  acid ;  for  these  acids  have  been  found  by  Furry  f 
to  be  incapable  of  liberating  iodine  regularly  from  a  mixture  of 
iodide  and  iodate.  Conditions  which  tend  to  increase  the  acidic 
activity  of  this  compound  are  high  concentrations  and  moderately 
low  tempera tures.f 

It  has  not  been  found  possible  under  any  conditions  to  rely 
upon  the  immediate  liberation  of  the  full  amount  of  iodine;  a 
certain  period  of  time  is  required  for  the  completion  of  the 
reaction.  When  the  solution  is  of  small  volume  and  saturated 
with  mannite,  the  reaction  goes  tp  the  end  most  quickly  — 

*  Am.  Jour.  Sci.,  [4],  viii,  127. 

t  Am.  Chem.  Jour.,  vi,  341. 

J  Magnanini,  Gaz.  Chim.,  xx,  428;  xxi,  134,  1016,  1017. 


BORON  '.  211 

sometimes  almost  immediately ;  but  if  the  solution  of  boric  acid 
is  too  concentrated  —  nearly  saturated  —  the  boric  acid  alone 
throws  out  some  iodine  from  the  iodide-iodate  mixture  added 
to  destroy  other  free  acid,  and  on  bleaching  with  thiosulphate 
a  starting  point  is  obtained  at  which  some  of  the  boric  acid 
has  already  entered  into  combination.  The  amount  of  iodine 
thus  liberated  by  the  boric  acid  is,  however,  not  large,  and  if, 
upon  the  addition  of  the  iodide  and  iodate,  the  iodine  thrown  out 
by  the  free  hydrochloric  acid  present  is  immediately  bleached  by 
thiosulphate  and  the  analysis  proceeded  with  from  this  as  the 
neutral  point,  even  in  concentrated  solutions  the  error  is  almost 
inappreciable.  If,  however,  considerable  time  intervenes  be- 
tween the  adding  of  the  iodide  and  iodate  and  the  determination 
of  the  neutral  point  by  thiosulphate,  iodine  equivalent  to  as 
much  as  several  milligrams  of  boric  acid  may  be  liberated. 
This  difficulty  was  not  met  with  in  those  experiments  in  which 
the  iodide  and  iodate  were  added  to  solutions  of  concentration 
like  that  of  the  standard  solution  used  (7.738  grm.  per  liter),  but 
in  an  attempt  to  estimate  the  boric  acid  in  colemanite,  where 
the  solution  was  kept  as  concentrated  as  possible,  hoping  in  this 
way  to  decrease  the  time  required  for  the  complete  liberation  of 
iodine,  low  values  were  obtained;  that  is,  a  false  starting  point 
was  used.  At  the  time  of  adding  the  iodide  and  iodate  the  vol- 
ume should  not  be  less  than  25  cm.3  for  each  decigram  of  boric 
anhydride  (B2Oi)  present,  and  should  not  be  much  greater  than 
two  or  three  times  that  amount.  At  lower  concentrations  of 
the  boric  acid,  even  though  the  liquid  be  saturated  with  mannite, 
the  necessary  time  of  standing  is  prolonged  and  the  effect  of  car- 
bon dioxide  upon  the  iodide  and  iodate  is  increased ;  for  carbon 
dioxide,  whether  derived  from  the  atmosphere  or  existing  dis- 
solved in  the  solution,  upon  standing  slowly  liberates  iodine. 
The  effect  of  carbon  dioxide  is,  however,  small,  and  in  the  time 
required  for  the  completion  of  the  process  has  never  been  found 
equivalent  to  more  than  a  single  drop  of  the  solution  of  thiosul- 
phate used.  Even  if  the  material  to  be  analyzed  contains 
carbonates,  after  acidifying  in  concentrated  solution  and  shaking 
vigorously  the  small  amount  of  uncombined  carbon  dioxide 
remaining  has  an  almost  inappreciable  effect  upon  the  results. 
The  length  of  time  required  for  the  liberation  of  the  theoretical 
amount  of  iodine  in  a  solution  of  25  cm.3  to  50  cm.3  to  each 


212  METHODS  IN  CHEMICAL  ANALYSIS 

o.i  grm.  of  boric  anhydride  is  20  to  45  minutes,  and  at  the 
end  of  45  minutes'  standing  in  a  solution  saturated  with  man- 
nite  the  reaction  may  be  considered  complete.  During  this 
period,  however,  it  is  well  to  keep  the  solution  cool  —  zero  tem- 
perature will  do  no  harm  —  and  to  insure  thorough  mixture  by 
occasional  shaking.  As  free  iodine  would  tend  to  escape  upon 
standing  unless  kept  in  a  closed  flask,  it  is  convenient,  immedi- 
ately after  the  addition  of  mannite,  to  treat  with  an  excess  of  the 
standard  solution  of  thiosulphate,  —  8  cm.3  or  iocm.3  more  than 
the  amount  required  to  bleach  the  iodine  liberated,  — and  at  the 
expiration  of  40  to  60  minutes  to  titrate  back  with  n/io  iodine. 
The  strength  of  the  thiosulphate  solution  found  most  convenient 
is  n/5,  while  the  use  of  iodine  of  one-half  this  strength,  n/io, 
diminishes  the  error  of  reading  correspondingly.  In  solutions 
of  the  volume  recommended  the  addition  of  starch  to  give  the 
indication  with  iodine  is  unnecessary,  since  a  single  drop  of  one- 
twentieth  normal  iodine  in  excess  is  sufficient  to  give  a  strong 
lemon  coloration. 

Procedure  in  Summary.  —  The  procedure  recommended  is  as 
follows :  The  borate  is  dissolved  in  as  small  volume  and  as  little 
hydrochloric  acid  as  possible,  with  shaking  to  remove  free  carbon 
dioxide  and  adjustment  of  volume  so  that  at  the  time  of  adding 
potassium  iodide  and  iodate  there  shall  be  approximately  25  cm.3- 
50  cm.3  of  solution  for  each  decigram  of  boric  anhydride  present. 
The  greater  part  of  the  excess  of  hydrochloric  acid  in  the  solution 
is  destroyed  by  sodium  hydroxide,  with  the  use  of  litmus  paper 
as  an  indicator,  leaving  the  solution  distinctly  acid  in  reaction. 
Potassium  iodide  (3  cm.3-5  cm.3  of  a  40  per  cent  solution)  and 
iodate  (5  cm.3-io  cm.3  of  a  5  per  cent  solution)  are  added  in  excess 
of  that  required  to  liberate  iodine  in  an  amount  corresponding  to 
the  hydrochloric  acid  and  the  boric  acid  present.  The  iodine 
liberated  by  the  free  hydrochloric  acid  is  bleached  by  a  small 
amount  of  a  strong  solution  of  thiosulphate,  and,  after  agitating 
to  insure  thorough  mixture,  iodine  is  added  to  faint  coloration. 
Sufficient  mannite  is  now  used  to  saturate  the  solution  —  about 
10  grm.-i5  grm.  for  a  volume  of  50  cm.3  —  and  sodium  thiosul- 
phate is  added  in  standard  solution  8  cm.3-io  cm.3  in  excess  of 
that  required  to  bleach  the  iodine  immediately  thrown  out  by 
the  mannite.  The  solution  is  again  brought  to  saturation,  if  nec- 
essary, by  mannite,  and,  after  standing  in  a  cool  place  for  40-60 


BORON 


213 


minutes,  titrated  with  decinormal  iodine  to  determine  the  excess 
of  thiosulphate  present. 

The  results  of  experiments  upon  pure  boric  acid,  crude  calcium 
borate  and  crystallized  colemanite  are  given  in  the  accompany- 
ing tables. 

Pure  Boric  Acid. 


B203* 

taken. 

cm  .3 

Thiosul- 
phate f 
taken. 

cm.3 

lodinej 
taken. 

cm.3 

Time  of 
standing. 

min. 

Volume. 
cm.8 

B203 

taken. 

grm. 

B2O3  found, 
grm. 

Error, 
grm. 

28.00 

32.00 

1.88 

30 

28 

0.2165 

0.2168 

+0.0003 

27.03 

32.00 

4-37 

27 

27 

0.2090 

o.  2081 

—  0.0009 

27.02 

31-97 

4.04 

60 

27 

0.2089 

o  .  2090 

+0  .  OOOI 

27.06 

32.04 

3.88 

60 

50-60 

0.2093 

O.  2IOI 

+0.0008 

27.02 

32.02 

4.40 

60 

50-60 

0.2089 

o.  2081 

—  0.0008 

27.04 

31.72 

3-39 

60 

50-60 

0.2091 

o  .  2096 

+0.0005 

27.01 

3J-53 

2.88 

1  2O 

50-60 

0.2089 

O.  2IOO 

+O.OOII 

26.05 

31.01 

4.01 

1  80 

50-60 

0.2014 

0.2025 

+O.OOII 

27.00 

31.00 

2.  12 

3° 

50-60 

0.2088 

o  .  2089 

+0.0001 

27.00 

32.00 

4-05 

30 

50-60 

0.2088 

O.2O92 

+0.0004 

26.01 

32.02 

6.  20 

30 

50-60 

O.  2OII 

0.2018 

+o  .  0007 

27.03 

31.01 

2.21 

48 

50-60 

O.2O9O 

0.2087 

—0.0003 

27.05 

31.89 

3-8l 

45 

50-60 

O.2092 

0.2093 

+O  .  OOOI 

26.07 

31.02 

4.14 

40 

50-60 

o.  2016 

O.  2O2O 

+0.0004 

27.00 

32.04 

4-30 

40 

60 

0.2088 

o.  2086 

—  O.OOO2 

*  7.773  grm.  per  liter. 


t  0.198  normal. 
Calcium  Borate. 


J  0.0996  normal. 


Mineral. 

Thiosul- 
phate taken. 

Iodine 
taken. 

Time  of 
standing. 

Volume  of 
solutions. 

B2O3  found. 

Per  cent. 

grm. 

cm.3 

cm.3 

min. 

cm.3 

grm. 

0.4015 
0.4010 

35-05 
35-34 

4-75 
5-23 

60 
1  2O 

40 
45 

O.  2280 
0.2283 

56.92 
56.94 

Colemanite. 

o  .  4002 

32 

.00 

5-50 

90 

50 

o 

2043 

51.04 

0.2513 

32 

.01 

7.36 

60 

40 

o 

1279 

50.91 

0.4007 

33 

•03 

7.72 

So 

65 

o 

2036 

50.81 

These  results  show  little  variation,  and  in  the  case  of  cole- 
manite correspond  closely  to  the  theory  50.97  per  cent.  The 
process  is  convenient,  generally  applicable,  and  accurate  within 
the  ordinary  limits  of  analysis. 


214  METHODS  IN  CHEMICAL  ANALYSIS 


ALUMINIUM. 

The  Determination  of  Aluminium  by  Precipitation  with  Ether- 
Hydrochloric  Acid. 

Crude  aluminium  chloride  may  be  freed  from  every  trace  of  a 
ferric  salt  by  dissolving  it  in  the  least  possible  amount  of  water, 
saturating  the  cooled  solution  with  gaseous  hydrochloric  acid, 
filtering  upon  asbestos  in  a  filtering  crucible  or  cone,  and  wash- 
ing the  crystalline  precipitate  with  the  strongest  hydrochloric 
acid.  Prepared  in  this  way  the  salt  gives  no  trace  of  color  when 
dissolved  in  water  and  tested  with  potassium  sulphocyanate, 
but  the  degree  of  insolubility  is  not  sufficient  for  the  pur- 
poses of  analysis.  Gooch  and  Havens*  have  found  that  in  a 
mixture  of  hydrochloric  acid  of  highest  concentration  and  ether 
in  equal  parts  the  solubility  of  aluminium  chloride  amounts 
approximately  to  5  parts  of  the  hydrous  salt,  A1C13.6H2O,  cor- 
responding to  I  part  of  the  oxide,  A^Oa,  in  125,000  parts  of  the 
mixture. 

Pure  aqueous  hydrochloric  acid  of  full  strength  mixes  per- 
fectly with  its  own  volume  of  anhydrous  ether,  but  the  addition 
to  this  mixture  of  any  very  considerable  amount  of  a  solution 
of  ferric  chloride  in  strong  hydrochloric  acid  occasions  the  sepa- 
ration of  a  greenish  etherial  solution  of  the  ferric  salt  upon  the 
surface  of  the  acid.  The  addition  of  more  aqueous  acid  does  not 
change  the  conditions  essentially,  but  more  ether  renders  the 
acid  and  the  oily  solution  completely  miscible.  For  the  separa- 
tion of  insoluble  aluminium  chloride  from  certain  small  amounts 
of  soluble  ferric  chloride,  the  mixture  of  the  strongest  aqueous 
hydrochloric  acid  and  ether  in  equal  parts  serves  the  purpose 
excellently;  when  larger  amounts  of  ferric  chloride  are  to  be 
dissolved,  ether  must  be  added  proportionately  in  order  to  pre- 
vent the  separation  of  the  etherial  solution  of  ferric  chloride  from 
the  rest  of  the  liquid. 

separation  of  The  quantitative  procedure  as  finally  developed 
Aluminium  by  Gooch  and  Havens  for  the  separation  of  alumin- 
ium from  iron,  and  the  determination  of  aluminium 
as  the  oxide,  is  as  follows :  The  concentrated  aqueous  solution  of 
the  salts  is  mixed  with  a  suitable  volume  of  strongest  aqueous 

*  F.  A.  Gooch  and  F.  S.  Havens,  Am.  Jour.  Sci.,  [4],  ii,  416. 


ALUMINIUM 


hydrochloric  acid  —  enough  to  make  the  entire  volume  approxi- 
mately 15  cm.3  to  25  cm.3.  This  mixture  is  saturated  with  gase- 
ous hydrochloric  acid  while  kept  cool  by  immersing  in  running 
water  the  receptacle  containing  it.  A  volume  of  ether  equal  to 
the  volume  of  the  liquid  is  introduced,  and  the  cooled  etherial 
mixture  is  treated  with  gaseous  hydrochloric  acid  to  saturation. 
The  precipitated  crystalline  chloride  is  collected  upon  asbestos 
in  a  perforated  crucible,  washed  with  a  previously  prepared  mix- 
ture of  hydrochloric  acid  and  ether  carefully  saturated  at  15°, 
dried  a  half-hour  at  150°,  covered  with  a  layer  of  pure  mercury 
oxide*  (about  I  grm.),  and  ignited  carefully  under  a  good  ven- 
tilating flue,  finally  with  the  blast  lamp. 

The  gaseous  hydrochloric  acid  is  most  conveniently  produced 
in  regulated  current  by  treating  massive  ammonium  chloride 
with  strong  sulphuric  acid  in  the  Kipp  generator.  A  platinum 
dish  hung  in  an  inverted  bell  jar,  provided  with  inlet  and  outlet 
tubes  through  which  the  current  of  water  for  cooling  is  passed, 
makes  a  convenient  container  for  the  solution  to  be  saturated 
with  the  gas.  The  filtration  is  made  upon  asbestos  in  a  perfo- 
rated crucible.  The  filtrate  and  washings  are  caught  directly  in 
a  crucible  (placed  under  the  bell  jar  of  the  filter  pump)  in  which 
the  subsequent  evaporation  is  to  be  effected.  The  heating  of 
the  strongly  acid  solution  must  be  gradual  and  conducted  with 
care  to  prevent  mechanical  loss  by  a  too  violent  evolution  of  the 
gaseous  acid. 

Results  are  given  in  the  table. 


Precipitation  of  Aluminium  Chloride  in  Presence  of  Ferric  Chloride. 


A12O3  taken  in 
solution  as  the 
chloride. 

Al2Os  found  by 
ignition  with  HgO. 

FejOs  present  as 
chloride. 

Final  volume. 

Error. 

grm. 

grm. 

grm. 

cm.s 

grm. 

0.0761 

0.0758 

25 

—  0.0063 

0.0761 

0.0754 

25 

—  0.0007 

0.0761 

0.0751 

.... 

25 

—  o.ooio 

0.0761 

0.0757 

0.15 

25-30 

—0.0004 

0.0761 

0.0756 

0.15 

25-30 

—0.0005 

0.0761 

0-0755 

0.15 

25-30 

—0.0006 

0.0761 

0.0755 

0.15 

25-30 

—  0.0006 

*  Loc.  cit.,  p.  419. 


2l6 


METHODS  IN  CHEMICAL  ANALYSIS 


Determination  By  similar  procedure  Havens*  has  effected  the  sep- 
of  Aluminium  aration  of  aluminium  from  beryllium.  The  beryl- 
Beryllium.  jjum  ma^  ^  recovere(j  jn  ^6  filtrate  from  the  alu- 
minium chloride  by  precipitation  as  hydroxide  with  ammonia 
after  nearly  complete  evaporation  of  the  acid,  and  weighed  as 
oxide  after  ignition.  Or,  the  filtrate  may  be  evaporated  just 
to  dryness  on  a  radiator,  care  being  taken  not  to  heat  to  the 
volatilizing  point  of  the  beryllium  chloride  ;  a  few  drops  of  strong 
nitric  acid  added;  the  liquid  evaporated,  best  with  a  current  of 
air  playing  on  the  surface;  and  the  residue  heated  gently  at 
first,  to  break  up  the  nitrate,  and  finally  over  the  blast  lamp. 
Results  of  this  procedure  are  given  below. 


Aluminium  and  Beryllium. 


Al2Oj  taken  in 
solution  as  the 

A1203. 

Error. 

Final 
volume. 

BeO  taken  in 
solution  as  the 

BeO  found. 

Error. 

chloride. 

chloride. 

grm. 

grm. 

grm. 

cm.8 

grm. 

grm. 

grm. 

0.1059 

0.1058 

—  O.OOOI 

12 

0.0198 

O  .  0204 

+0.0006* 

O.IOS3 

o.  1044 

—  0.0009 

12 

0.0194 

0.0196 

-fo.0002* 

0.1065 

o.  1059 

—  0.0006 

12 

0.0197 

0.0205 

+0.0008* 

0.1068 

o.  1060 

—  0.0008 

12 

0.0199 

0.0207 

+0.0008* 

0.1049 

o.  1047 

—  O.OOO2 

12 

0.0198 

0.0208 

+O.OOIO* 

o.  1060 

,  0.1057 

—0.0003 

12 

0.0977 

O  .  0969 

-0.0008* 

o.  1064 

0.1063 

—  O.OOOI 

12 

0.1085 

o.  1084 

—  O.OOOI* 

o.  1046 

o.  1038 

—0.0008 

30 

0.1083 

o.  1087 

+0.0004* 

0.1051 

o.  1048 

—0.0003 

30 

o.  1071 

0.1078 

+0.0007* 

0.1076 

0.1075 

—  O.OOOI 

30 

0.1086 

0.1094 

+o.ooo8f 

*  By  the  evaporation  process, 
t  By  the  precipitation  process. 


Determination  Havensf  effected  similarly  the  separation  of  alu- 
of  Aluminium  minium  from  zinc,  the  zinc  chloride  in  the  filtrate  be- 
andZmc.  -^  converteci  to  the  oxide  by  evaporation,  repeated 

treatments  with  nitric  acid  in  porcelain,  and  ignition.  On  ac- 
count of  the  danger  to  platinum,  the  evaporations  and  treat- 
ments with  nitric  acid  are  made  in  porcelain.  In  these  processes, 
however,  the  porcelain  is  somewhat  attacked,  and  correction 
must  be  made  for  a  slight  contamination  of  the  residual  oxide. 
Results  are  given  on  the  following  page. 

*  Franke  S.  Havens,  Am.  Jour.  Sci.  [4],  iv,  HI. 
f  Franke  Stuart  Havens,  Am.  Jour.  Sci.,  [4],  vi,  45. 


ALUMINIUM 


217 


Aluminium  and  Zinc. 


A120, 

taken  as 
the 

A1203 
found. 

Error. 

ZnO 

taken. 

ZnO 

found. 

Error. 

Error 
corrected.* 

Final 
volume. 

chloride. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

cm.3 

o  0^62 

O   OOOO 

O    IIIO 

0.0580 

0.0577 

—  0.0003 

0.1034 

0.0572 

0.0572 

O  .  OOOO 

o.  1014 

0.1027 

+0.0013 

—  0.0007 

12 

0.0563 

0.0550 

-0.0013 

0.1026 

o.  1038 

+0.0012 

—  0.0008 

16 

0.0577 

0.0576 

—  O.OOOI 

O.IOOO 

O.IOI4 

+O.OOI4 

—  O.OOO6 

16 

o-°559 

0.0558 

—  O.OOOI 

O.IO2O 

0.1035 

+  O.OOI5 

—  0.0005 

16 

0.0563 

0.0556 

—0.0007 

0.2024 

o  .  2046 

+O.OO22 

+O  .  OOO2 

20 

O.IIII 

o.  1107 

—0.0004 

0.2092 

0.2116 

+O.OO24 

+0  .  0004 

20 

*  Corrected  by  the  amount  of  material  found  in  the  evaporation  of  similar  amounts  of  the  strong 
acids  in  porcelain. 

Determination  The  process  applies  also  to  the  separation  of  alu- 
of  Aluminium  minium  from  copper,  as  shown  by  the  results  given.* 
and  copper.  The  copper  in  the  filtrates  of  these  determinations 
was  converted  to  the  oxide  through  the  sulphate  because  this 
operation  may  be  conducted  safely  in  platinum. 

Aluminium  and  Copper. 


A12O3  taken  as 
chloride. 

grm. 

A12O3  found, 
grm. 

Error, 
grm. 

CuO  taken, 
grm. 

CuO  found, 
grm. 

Error, 
grm. 

0.0437 

0  .  043  2 

—  O  0005 

0.0359 

O.O35Q 

O   OOOO 

O  O34<? 

o  0340 

—  o  0005 

0.0558 
0.0538 
0.0566 
0.0577 

0.0545 
0.0536 
0.0562 
0.0575 

—  0.0013 
—  O.OOO2 
—  O.0004 
—  O.OOO2 

0.0319 
0-0343 
0.0337 
0.0651 

0.0324 
0.0356 
0.0349 

o  .  0644 

+O.OOO5 
+0.0013 
+0.0012 
—  0.0007 

Separation  of 
Aluminium 
from  Mercury 
and  Bismuth. 


Aluminium  may  be  similarly  separated  from  mer- 
cury and  bismuth,  as  shown  by  the  results  given 
below.f 

Aluminium,  Mercury  and  Bismuth. 


A12O3  taken  as 
chloride. 

A12O3  found. 

Error. 

HgCl2  taken. 

Bi2O3  taken. 

grm. 

grm. 

grm. 

grm. 

grm. 

0.0570 
0  0548 
0.0565 
0.0576 

0.0574 
0.0557 
0.0571 
0.0577 

+0.0004 
+0  .  0009 
+0  .  0006 
+0  .  0001 

O.I 
O.  I 

O.I 
O.2 

*  Havens,  loc.  cit. 


t  Havens,  loc.  cit. 


218 


METHODS  IN  CHEMICAL  ANALYSIS 


LANTHANUM. 
The  Estimation  of  Lanthanum  Precipitated  as  the  Oxalate. 

Many  years  ago  Stolba*  stated  that  cerium,  lanthanum  and 
didymium  may  be  estimated  by  treating  their  oxalates  with 
potassium  permanganate  in  the  presence  of  sulphuric  acid,  but 
gave  no  experimental  evidence  in  the  form  of  analytical  results. 
Later  this  statement  was  confirmed  by  Browning  and  Lynch,  f 
and  it  was  shown  that  cerium  may  be  estimated  by  precipitating 
cerium  oxalate  with  a  definite  amount  of  a  standard  solution  of 
ammonium  oxalate  used  in  excess,  decomposing  the  precipi- 
tated cerium  oxalate  by  dilute  sulphuric  acid  and  estimating 
the  oxalate  by  permanganate.  The  ammonium  oxalate  in  ex- 
cess of  the  amount  required  for  the  precipitation  was  also 
estimated  by  permanganate,  and  by  this  process  the  results  were 

checked. 

.    .  Lanthanum  Precipitated  as  the  Oxalate. 


LajO3  taken.* 
grin. 

La2O3  found,  t 

Average, 
grm. 

Error, 
grm. 

Precipitate, 
grm. 

Filtrate, 
grm. 

0.0148 

0.0152 

0.0144 

0.0148 

O  .  oooo 

0.0148 

0.0149 

0.0139 

0.0144 

—  0.0004 

0.0296 

o  .  0302 

0.0291 

0.0296 

o.oooo 

0.0296 

o  .  0302 

0.0293 

0.0297 

+O.OOOI 

0.0592 

0.0599 

0.0586 

0-0593 

+O.OOOI 

0.0592 

0.0598 

0.0585 

0.0592 

0.0000 

0.1184 

o.  1191 

0.1179 

o.  1185 

+O  .  OOOI 

0.1184 

o.  1191 

o.  1182 

o.  1187 

+0  .  0003 

0.2368 

0.2376 

0.2362 

0.2369 

-f-o.oooi 

0.0148 

0.0149 

0.0145 

0.0147 

—  O.OOOI 

0.0148 

0.0150 

0.0147 

0.0148 

o.oooo 

0.0296 

0.0298 

o  .  0293 

0.0295 

—  O.OOOI 

0.0592 

0.0596 

0.0589 

0.0593 

+0.0001 

0.1184 

o.  1190 

o.  1182 

0.1186 

+O.OOO2 

0.1036 

o.  1040 

o.  1029 

0.1035 

—  O.OOOI 

*  Taken  as  ammonium  lanthanum  nitrate. 


t  La =138.9. 


The  conditions  under  which  lanthanum  may  be  estimated  as 
the  oxalate  have  been  studied  by  Drushel  |  and  a  process  of 
treatment  recommended.  The  procedure  is  as  follows:  • 

*  Sitzungsber.  d.  kgl.  bohm.  Gesellsch.  d.Wissenschaften,  v,  4,  Juli,  1879. 

t  See  page  248. 

t  W.  A.  Drushel,  Am.  Jour.  Sci.,  [4],  xxiv,  197. 


THALLIUM  219 

From  a  neutral  solution  of  a  lanthanum  salt,  conveniently  the 
nitrate,  the  oxalate  is  precipitated  by  a  measured  amount  of 
standard  n/io  oxalic  acid,  or  ammonium  oxalate,  after  the  addi- 
tion of  a  few  drops  of  acetic  acid.  The  precipitate  is  thoroughly 
stirred,  allowed  to  settle,  and  filtered  off  in  a  perforated  crucible 
fitted  with  an  asbestos  felt.  After  thoroughly  washing  with 
water  the  crucible  and  precipitate  are  placed  in  a  beaker  with 
100  cm.3  to  300  cm.3  of  water  and  10  cm.3  to  30  cm.3  of  [i  :  3] 
sulphuric  acid.  The  contents  of  the  beaker  are  heated  nearly 
to  boiling,  and  at  once  titrated  to  color  with  standard  potas- 
sium permanganate.  The  filtrate  is  similarly  titrated  as  a  check 
on  the  titration  of  the  precipitate.  The  lanthanum  is  calcu- 
lated as  La2O3  from  the  two  titrations.  The  mean  of  the  two 
closely  agreeing  values  thus  obtained  is  taken.  Results  are 
shown  on  the  preceding  page. 

THALLIUM. 

The  Determination  of  Thallium  as  the  Acid  Sulphate  and  as  the 
Neutral  Sulphate. 

Crookes*  has  shown  that  the  salt  obtained  by  heating  thallous 
chloride  with  sulphuric  acid  until  the  excess  of  the  latter  is 
expelled,  and  then  raising  the  heat  to  redness,  has  the  consti- 
tution of  a  neutral  sulphate  and  sustains  no  further  appreciable 
loss  of  weight  on  heating,  and  he  has  suggested  the  possibility  of 
applying  this  treatment  in  the  estimation  of  thallium.  Castanjenf 
essentially  confirms  these  statements  in  regard  to  the  sulphate, 
adding  the  observation  that  the  salt  tends  to  lose  acid  on  strong 
ignition  in  air.  Experiments  by  Browning  J  show,  however,  that 
with  proper  precautions  thallium  may  be  weighed  either  as  the 
neutral  sulphate  or  as  the  acid  sulphate. 

The  procedure  is  as  follows:  To  the  solution  of  the  thallium 
salt  of  a  volatile  acid,  taken  in  a  crucible,  sulphuric  acid  is  added, 
and  the  water  is  removed  as  far  as  possible  by  evaporation  oa 
the  steam  bath.  The  crucible  is  then  placed  within  a  conical 
iron  radiator  and  heated  at  temperatures  ranging  between  220° 
and  240°  until  fuming  ceases  and  the  weights  after  half-hour 

*  Chem.  News,  viii,  243. 

t  Jour,  prakt.  Chem.,  cii,  131. 

J  Philip  E.  Browning,  Am.  Jour.  Sci.,  [4],  ix,  137. 


220 


METHODS  IN  CHEMICAL  ANALYSIS 


periods  of  heating  remain  constant.     These  weights  give  the 
weight  of  acid  thallium  sulphate. 

Upon  further  heating  of  the  crucible  and  contents  over  a 
free  flame  to  low  redness  the  weight  again  becomes  constant, 
after  considerable  evolution  of  fumes,  and  shows  a  condition 
closely  approximating  that  of  neutral  thallium  sulphate.  Ex- 
perimental results  of  this  treatment  of  thallium  nitrate  are 
given  in  the  table. 

Conversion  of  Thallium  Nitrate  to  Thallium  Sulphates. 


T1HSO4 
calculated. 

grm. 

T1HS04 
found. 

grm. 

Error, 
grm. 

T12S04 
calculated. 

grm. 

T12SO< 
found. 

grm. 

Error. 

gnu. 

o.  1605 
o.  1611 
0.1608 
0.1612 
o.  1602 
0.1608 

0.1596 
0.1608 
0.1608 
o.  1600 
0.1596 
o  1^06 

—  0.0009 
—  0.0003 
O  .  OOOO 
—  O.OOI2 
—  O.OOO6 
—  O   OOI2 

0.1344 
0.1349 
0.1347 
0.1350 
0.1341 

0.1346 
0.1346 
0.1352 
0.1346 
0.1346 

+  0   0002 

—  o  0003 
+o  0005 

—  O.O004 
+0.0005 

o.  1617 

o  1604 

—  o  0013 

0.1608 
0.1609 

0.1592 
0.1590 

—  0.0016 
—0.0019 

0.1347 
0.1348 

0.1358 
0.1346 

+O.OOII 
—  O.OO02 

The  Gravimetric  Estimation  of  Thallium  Precipitated  as  Thallic 
Hydroxide  by  Potassium  Ferricyanide  and  Potas- 
sium Hydroxide. 

The  precipitation  of  thallic  hydroxide  *  by  the  action  of  po- 
tassium ferricyanide  and  potassium-  hydroxide  has  been  recom- 
mended by  Browning  and  Palmer, f  as  satisfactory  means  for 
the  separation  of  thallium  in  easily  determinable  form  according 
to  the  following  procedure. 

To  the  solution  of  the  thallous  salt  in  100  cm.3  of  water  are 
added  a  solution  of  potassium  ferricyanide,  in  excess,  and  potas- 
sium hydroxide  to  complete  precipitation  of  the  brown  thallic 
hydroxide.  The  precipitate  is  filtered  off  on  asbestos  in  the 
perforated  crucible,  thoroughly  washed,  best  with  hot  water, 
and  dried  over  a  low  flame,  at  about  200°,  to  constant  weight. 
Test  results  are  given  in  the  table. 

*  See  page  223. 

t  Philip  E.  Browning  and  Howard  E.  Palmer,  Am.  Jour.  Sci.,  [4],  xxvii.  380. 


THALLIUM 


221 


Precipitation  as  Thallic  Hydroxide. 


T12O3  taken  as 
thallous  nitrate, 
giro. 

T12O3  found, 
grin. 

Error, 
gnn. 

0.1305 

0.1309 

+o  .  0004 

0.1305 

0.1314 

+0.0009 

0.1305 

o.  1308 

-j-o  .  0003 

0.0870 

0.0872 

+O.OOO2 

0.1740 

0.1741 

+O.OOOI 

o  .  i  740 

0.1739 

—  O.OOOI 

o  .  i  740 

0.1742 

+O.OOO2 

0.1305 

0.1307 

+O.OOO2 

0.1305 

0.1309 

+0.0004 

0.1305 

o  .  1308 

+0.0003 

0.0870 

0.0872 

+O.OOO2 

0.0870 

0.0874 

+O.OOO4 

The  Gravimetric  Estimation  of  Thallium  as  the  Chromate. 

Crookes  has  shown  that  thallium  chromate  precipitated  by  the 
addition  of  potassium  dichromate  to  an  alkaline  solution  of  a 
thallous  salt  has  the  constitution  of  a  neutral  salt,  is  very  in- 
soluble in  water,  — 100  parts  of  water  at  100°  dissolving  about 
0.2  parts  and  at  60°  about  0.03  parts,  —  and  may  be  used  to 
effect  a  rough  separation  of  thallium  from  cadmium.  Browning 
and  Hutchins  *  have  described  conditions  under  which  the  reac- 
tion gives  exact  quantitative  results.  To  the  solution  of  thallous 
nitrate,  heated  to  70°  or  80°,  are  added,  with  shaking,  ammonium 
hydroxide,  or,  better,  potassium  carbonate,  to  alkalinity,  and  an 
excess  of  potassium  dichromate  in  solution.  After  cooling  and 
settling,  the  thallous  chromate  is  filtered  off  on  asbestos  in  the 
perforated  crucible  and  dried  over  a  low  flame  to  constant 
weight.  If  the  precipitation  is  made  in  the  cold,  the  chromate 
does  not  form  well  but  remains  partly  in  finely  divided  condition 
and  difficult  to  filter.  The  addition  of  ammonium  nitrate  before 
precipitation  brings  about  a  better  condition  even  in  the  cold; 
but  the  best  results  are  obtained  by  precipitating  the  warm  solu- 
tion made  alkaline  with  potassium  carbonate.  Results  follow 
in  the  table. 


*  Philip  E.  Browning  and  George  P.  Hutchins,  Am.  Jour.  Sci.,  [4],  viii,  460. 


222 


METHODS  IN  CHEMICAL  ANALYSIS 


Determination  of  Thallium  as  Thallous  Chr ornate. 


T1NO3  taken. 
Calculated  as  T12O. 

grni. 

Tl2CrO4  found. 
Calculated  as  T12O. 

gnu. 

Error. 
Calculated  as  T12O. 

grin. 

o  .  0796 

0.0791 

-0.0005 

0.0792 

0.0788 

—0.0004 

0.0792 

0.0786 

—0.0006 

0.1188 

0.1177 

—  O.OOII 

o.  1192 

O.II86 

—0.0006 

o.  1185 

o.  1178 

—0.0007 

o.  1190 

o.  1185 

—0.0005 

0.1189 

0.1183 

—0.0006 

0.1196 

O.2OOO 

+0.0004 

0.1196 

0.2005 

-j-o  .  0009 

0.1173 

O.II73 

o.oooo 

o.  1171 

O.II63 

—0.0008 

The  lodometric  Estimation  of  Thallium  by  Precipitation  with  Potas- 
sium Dichromate  and  Determination  of  the  Excess 
of  the  Precipitant. 

Browning  and  Hutchins*  have  shown  that  thallium  may  be 
estimated  by  precipitation  as  chromate  and  estimation  of  the 
excess  of  precipitant  left  over  from  a  definite  amount  taken. 

According  to  this  process,  a  solution  of  potassium  dichromate 
is  standardized  f  by  treatment  of  the  solution,  acidified  with 
sulphuric  acid,  by  a  definite  amount  of  standard  arsenite,  set- 
ting aside  until  the  change  from  yellow  to  bluish  green  indicates 
the  reduction  of  the  chromate,  making  alkaline  with  acid  potas- 
sium carbonate,  and  titrating  the  residual  arsenite  by  standard 
iodine.  To  the  solution  of  thallous  nitrate,  heated  to  70°  or  80° 
and  made  alkaline  with  acid  potassium  carbonate,  is  added  an 
excess  of  the  standardized  potassium  dichromate  and  the  precipi- 
tated thallous  chromate  is  filtered  off.  After  cooling,  the  filtrate 
containing  the  excess  of  alkali  chromate  is  acidified  with  sulphuric 
acid,  a  definite  amount  of  standard  arsenite  is  added,  and  the 
whole  allowed  to  stand  a  few  moments  until  the  change  from 
yellow  to  bluish  green  indicates  reduction  of  the  chromic  acid. 
Acid  potassium  carbonate  is  added  to  alkaline  reaction  and  the 
residual  arsenite  is  determined  by  titration  with  standard  iodine. 
The  difference  between  the  arsenite  taken  and  the  arsenite 

*  Philip  E.  Browning  and  George  P.  Hutchins,  Am.  Jour.  Sci.,  [4],  viii,  461. 
f  See  page  408. 


THALLIUM 


223 


found  measures  the  chromate  remaining  after  precipitating  the 
thallous  chromate,  and  the  difference  between  the  chromate  re- 
maining and  that  taken  is  the  chromate  equivalent  to  the  thal- 
lium. Results  of  this  procedure  are  given  in  the  table. 

Precipitation  of  Thallous  Chromate:  Determination  of  the  Excess  of  Precipitant. 


T12O  in  T1NO, 
taken. 

T120  in  Tl2Cr04 
found. 

Error. 

grm. 

grm. 

grm. 

0.1192 

o.  1198 

+0.0006 

0.  1189 

0  .  1  205 

+0.0016 

0.1196 

o.  1180 

—  0.0016 

o.  1196 

0.  IIQ2 

—0.0004 

O.II73 

o.  1182 

+o  .  0009 

o.  1171 

o.  1190 

+0.0019 

The  Estimation  of  Thallium  by  the  Action  of  Potassium  Ferricya- 

nide  in  Alkaline  Solution  and  of  Potassium  Permanganate 

in  Acid  Solution  upon  the  Ferrocyanide  Produced. 

The  oxidation  of  a  thallous  salt  with  precipitation  of  thallic 
hydroxide  by  the  action  of  potassium  ferricyanide  and  potassium 
hydroxide  and  the  subsequent  oxidation  of  the  potassium  ferro- 
cyanide  produced,  after  nitration  of  the  mixture,  are  the  essential 
steps  in  a  process  proposed  by  Browning  and  Palmer*  for  the 
volumetric  estimation  of  thallium. 

The  procedure  is  as  follows:  To  the  solution  of  the  thallous 
salt  in  about  100  cm.3  of  water  is  added  a  sufficient  excess  of  a 
solution  of  potassium  ferricyanide  and  potassium  hydroxide  to 
complete  precipitation  of  the  brown  thallic  hydroxide.  The 
precipitate  is  filtered  off  on  asbestos,  generally  without  settling, 
and  washed  thoroughly.  The  filtrate  is  acidified  with  sulphuric 
acid  and  titrated  with  standard  permanganate.  From  the  fol- 
lowing equations,  representing  the  reactions,  the  amount  of 
thallium  present  may  be  readily  calculated : 

T12O  +  4  KaFeCeNe  +  4  KOH  =  T12O3  +  4  K4FeC6N6  +  2  H2O, 
5  K4FeC6N6  +  KMnO4  +  4  H2SO4  = 

5  KsFeCeNe  +  3  K2SO4  +  MnSO4  +  4  H2O. 

It  is  necessary  to  apply  a  correction  for  the  amount  of  per- 
manganate used  to  give  the  first  tinge  of  pink  color  to  the  amounts 
*  Philip  E.  Browning  and  Howard  E.  Palmer,  Am.  Jour.  Sci.,  [4],  xxvii,  379. 


224 


METHODS  IN  CHEMICAL  ANALYSIS 


of   ferricyanide  used   in   the   determinations,   but  this  seldom 
exceeds  o.i  cm.3  of  the  permanganate. 

The  following  table  shows  the  results  obtained  with  different 
amounts  of  the  thallium  salt. 

Volumetric  Estimation  of  Thallium. 


T12O  taken  as  the 
nitrate. 

grm. 

T12O  found, 
grm. 

Error, 
grin. 

o  .  0809 

o  .  0809 

+O.OOOO 

O  .  0809 

O.o8o8 

—  O.OOOI 

0.0809 

0.0809 

o.oooo 

0.1213 

O.  1212 

—  O.OOOI 

0.1213 

o.  1216 

+0.0003 

0.1213 

o.  1218 

+0.0005 

0.1213 

o.  1218 

+0.0005 

0.1213 

O.I2I2 

—  O.OOOI 

0.1213 

0.1207 

—0.0006 

0.1618 

O.l6l4 

—0.0004 

0.1618 

O.l6l3 

—0.0005 

0.1618 

0.1616 

—  O.O002 

CHAPTER  VII. 


Carbon  Dioxide 
by  the  Action 
of  Acid. 


CARBON;  SILICON;  TITANIUM;  ZIRCONIUM;  CERIUM;  TIN;  LEAD. 

CARBON. 
The  Determination  of  Carbon  Dioxide  in  Carbonates  by  Loss. 

Expulsion  of          THE  determination  of  carbon  dioxide  in  carbonates, 
by  loss,  in  the  action  of  acid,  is  conveniently  made 
with  the   simple   apparatus  of   Kreider,   previously 
described  and  figured.* 

In  carrying  out  the  operation,  the  carbonate  is  weighed  and 
placed  in  the  bottom  of  the  test  tube,  A,  which  serves  as  the  re- 
action chamber.  The  acid,  to  a  volume  of  10  cm.3  to  15  cm.3,  is 
drawn  into  the  acid  chamber,  C,  and  held  there  in  the  manner 
described.  The  test  tube,  A,  is  slipped  over  B,  and  this  joint 
is  sealed  with  paraffin,  as  has  been  shown.  The  apparatus  is 
wiped,  placed  on  the  balance  and  weighed. 


Carbon  Dioxide  in  Carbonates. 


Taken, 
grin. 

Found, 
grm. 

Error, 
grin. 

O.2OOO 

0.0879 

—  O.OOOI 

O.  2OOO 

0.0878 

—  O.OOO2 

O.2OOO 

0.0879 

—  O.OOOI 

O.2OOO 

0.0879 

—  O.OOOI 

Calcium  carbonate  - 

0.5000 
0.5000 

0.2197 
0.2196 

—0.0003 
—0.0004 

0.5000 

0.2194 

—0.0006 

0.5000 

0.2198 

—  O.OOO2 

0.5000 

0.2197 

—0.0003 

0.5000 

0.2197 

—0.0003 

r 

0.5000 

0.1134 

—  O.OOII 

Barium  carbonate.           \ 

0.5000 

0.1137 

—0.0005 

O.5OOO 

0.1137 

—0.0005 

^ 

O.5OOO 

0.1136 

—0.0006 

Strontium  carbonate.  .  .  -I 

0.5000 
0.5000 
0.5000 

0.1485 
o.  1486 

0.1485 

—0.0004 
—0.0003 
—0.0004 

*  See  Fig.  I,  page  I. 
225 


226  METHODS  IN  CHEMICAL  ANALYSIS 

Upon  removing  the  cap  from  the  small  tube  in  C,  the  acid 
runs  from  C  into  A.  The  carbon  dioxide  liberated  is  forced  up- 
ward through  the  drying  column  of  calcium  chloride  and  escapes 
through  the  annular  space  between  B  and  C.  When  the  action 
ceases,  a  current  of  dry  air  is  forced  through  C,  to  remove  the 
carbon  dioxide,  the  cap  is  replaced,  and  the  apparatus  is  weighed. 
The  loss  of  weight  represents  the  carbon  dioxide. 

The  preceding  results  show  the  accuracy  which  may  be  ex- 
pected when  carbonates  are  treated  in  the  apparatus  with  dilute 
hydrochloric  acid. 

Expulsion  of  From  certain  carbonates,  like  those  of  magnesium, 

carbon  Dioxide  zinc,  and  cadmium,  carbon  dioxide  may  be  expelled 
by  simple  ignition  at  a  moderate  temperature,  leav- 
ing an  oxide  in  definite  and  weighable  condition.  In  the  case 
of  calcium  carbonate,  this  process  of  decomposition  is  completed 
only  at  the  high  heat  of  the  blast  lamp,  and  the  reaction,  being 
easily  reversible  in  the  atmosphere  containing  carbon  dioxide 
evolved  in  the  ignition  or  produced  in  the  source  of  heat,  may 
leave  the  oxide  not  quite  pure.  Strontium  carbonate  and  barium 
carbonate  are  not  entirely  broken  up  by  simple  ignition  under 
the  conditions  ordinarily  available  in  analysis ;  nor  are  the  alkali 
carbonates.  For  the  decomposition  of  refractory  carbonates  it 
is  customary  to  make  use  of  a  suitable  flux  which,  by  combining 
with  the  oxide,  will  aid  in  the  expulsion  of  the  carbon  dioxide. 
Anhydrous  borax,*  silicon  dioxide, f  potassium  dichromate,J  and 
recently,  sodium  metaphosphate,§  have  been  thus  used  in  the 
analysis  of  carbonates,  and  they  are  applicable  similarly  to  the 
determination  of  nitrogen  pentoxide  in  nitrates  which  leave 
definite  oxides  on  ignition.  Such  fluxes,  moreover,  serve  the 
very  essential  end  of  conserving  the  residual  oxides  in  definite 
and  stable  form  for  weighing  under  the  ordinary  atmospheric 
condition  of  the  balance  room.  Of  those  mentioned,  the  first 
two,  borax  and  silicon  dioxide,  require  prolonged  ignition  to  bring 
them  to  constant  weight  before  making  use  of  them  to  react  with 
the  carbonate  or  nitrate;  and  generally  they  yield  in  the  fusion 

*  Fresenius,    Zeit.  anal.  Chem.,  i,  181. 

t  Rose,  Ann.  Phys.,cxvi,  635;  Fresenius,  Zeit.  anal.  Chem.,  i.  184;  Richards 
and  Archibald,  Proc.  Am.  Acad.,  xxxviii,  443. 

t  Rose,  Ann.  Phys.,  cxvi,  131;  Fresenius,  Zeit.  anal.  Chem.,  i,  183. 

§  Lutz  and  Tschischikof,  Chem.  Zentralblatt,  1905,  i,  564;  Bottger,  Zeit. 
anal.  Chem.,  xlix,  487. 


CARBON  227 

process  a  pasty  magma,  so  that  prolonged  heat  at  a  high  tem- 
perature is  necessary  to  the  complete  expulsion  of  the  gaseous 
product.  Sodium  metaphosphate,  though  more  fluid  in  fusion, 
also  demands  prolonged  care  in  the  preparation.  Potassium 
dichromate  is  too  easily  decomposed  with  loss  of  oxygen  to  be 
employed  in  exact  processes  demanding  long-continued  fusion  or 
heating  to  temperatures  much  above  its  fusing  point.  These 
are  points  which  have  been  sufficiently  emphasized  in  the  work 
to  which  reference  has  been  made. 

It  has  been  shown  by  Gooch  and  Kuzirian*  that  in  sodium 
paratungstate,  of  composition  corresponding  approximately  to 
the  formulae  5Na2O.i2WO3,  or  NaioW^Ou,  we  have  material  very 
easily  prepared,  stable  in  fusion,  and  well  suited  for  use  as  a  flux 
in  the  rapid  determination  of  the  loss  of  carbonates  and  nitrates 
on  ignition.  This  sodium  paratungstate  is  prepared  by  dehy- 
drating and  fusing  over  the  blast  lamp  a  known  weight  of  normal 
sodium  tungstate,  Na2WO4.2H2O,  adding  an  equal  weight  of 
tungsten  trioxide,  WO3  (previously  ignited  with  care  to  remove 
all  ammonia  and  to  insure  complete  oxidation),  and  heating  to 
clear  fusion.  The  cooled  mass,  which  is  very  easily  pulverized, 
is  ground  in  a  mortar  and  bottled.  From  this  material,  kept  in 
a  desiccator  over  sulphuric  acid  (though  not  more  than  ordina- 
rily hygroscopic),  portions  are  weighed  for  the  analytical  deter- 
minations. Approximately  half  the  weight  of  the  paratungstate 
is  tungsten  trioxide  (molecular  weight  232),  and  this  should  be 
capable  of  expelling  carbon  dioxide  (molecular  weight  44)  to  the 
amount  of  one-fifth  its  own  weight.  In  practice,  the  weights 
of  paratungstate  used  should  exceed  ten  times  the  weight  of  car- 
bon dioxide.  It  is  best  to  weigh  a  platinum  crucible,  introduce 
the  dried  carbonate  and  weigh  again,  add  a  suitable  amount  of 
the  prepared  sodium  paratungstate,  stir  carefully  with  a  plati- 
num wire  with  care  to  avoid  mechanical  loss,  and  weigh  again. 
The  crucible  is  then  heated  over  a  Bunsen  burner,  first  at  very 
low  heat  and  then  to  fusion  of  the  mixture  for  five  minutes,  cooled 
in  a  desiccator  over  sulphuric  acid,  weighed,  and  reignited  to 
test  the  constancy  of  weight.  The  constant  weight  is  usually 
got  in  the  first  ignition.  In  the  following  tables  are  given  the 
results  of  the  estimation  of  carbon  dioxide  in  calcite,  and  in  the 
precipitated  carbonates  of  strontium  and  barium. 

*  F.  A.  Gooch  and  S.  B.  Kuzirian,  Am.  Jour.  Sci.,  [4],  xxxi,  497. 


228 


METHODS  IN  CHEMICAL  ANALYSIS 


Analysis  of  Calcium  Carbonate  (Calcite). 


CaCO3  taken. 

Na10W12041 
taken. 

Loss  on  ignition. 

Theory  for  CO8. 

Error. 

grrn. 

grm. 

grm. 

gnu. 

grm. 

0.5000 

2-5 

0.2195 

0.2198 

—  0.0003 

0.5000 

2-5 

0.2206 

0.2198 

—  0.0008 

0.5000 

2.5 

O.22OO 

0.2198 

—  O.OOO2 

0.5000 

2.5 

0.2203 

0.2198 

—  0.0005 

0.5000 

2.5 

O.22OO 

0.2198 

—  0.0002 

0.5000 

2.5 

O.22O4 

0.2198 

—  O.OO06 

0.5000 

2-5 

0.2IQO 

0.2198 

—  O.OO02 

0.5000 

2.5 

O.  22OO 

0.2198 

—  O.OOO2 

Analysis  of  Specially  Precipitated  Strontium  Carbonate.* 


SrCO3  taken. 

Na10W12041 
taken  (approx.). 

Loss  on  ignition. 

Theory  for  CO2. 

Error. 

grm. 

grm. 

grm. 

grm. 

grin. 

0.5000 
0.5000 
0.5000 

2-5 
2-5 
2-5 

0.1488 
0.1494 
0.1494 

0.1490 
0.1490 
0.1490 

—  0.0002 

+o  .  0004 
+o  .  0004 

0.5000 
0.5000 
0.5000 

2-5 
2-5 
2-5 

0.1490 
o.  1496 

0.1486 

0.141,0 
0.1490 
0.1490 

o.oooo 
+0.0006 
—0.0004 

Analysis  of  Specially  Precipitated  Barium  Carbonate.* 


BaCO3  taken. 

NauWtfO,, 

taken  (approx.). 

Loss  on  ignition. 

Theory  for  CO2. 

Error. 

grm. 

grm. 

grin. 

grm. 

grin. 

0.5000 

2-5 

O.  II2O 

0.1115 

+0.0005 

0.5000 

2.5 

O.II25 

o.  1115 

+0.0010 

0.5000 

2.5 

O.IIO9 

0.1115 

—  0.0004 

0.5000 

2.5 

O.III3 

0.1115 

—  O.OOO2 

0.5000 

2.5 

O.II23 

0.1115 

—  O.OOOS 

*  Prepared  from  the  pure  chloride  by  partially  precipitating  and  washing  with  hydrochloric 
acid,  dissolving  the  precipitated  chloride  in  water,  adding  the  solution  drop  by  drop  to  a  hot  satu- 
rated solution  of  ammonium  carbonate,  washing  the  precipitated  carbonate  and  drying  below  red 
heat. 


The  Precipitation  and  Gravimetric  Determination  of  Carbon 

Dioxide. 

The  estimation  of  carbon  dioxide  in  carbonates,  by  liberation 
of  that  gas  and  absorption  in  weighed  potash  bulbs,  demands 
the  careful  observance  of  precautions  and  the  expenditure  of 
much  time  and  attention.  According  to  a  method  proposed  by 


CARBON 


229 


Gooch  and  Phelps,*  the  rapid  absorption  of  carbon  dioxide 
evolved  by  the  action  of  acids  on  carbonates  is  effected  by 
barium  hydroxide  contained  in  a  specially  devised  apparatus; 
the  precipitated  barium  carbonate  is  filtered  off  and  washed 
under  a  protecting  layer  of  xylene;  the  washed  carbonate  on 
the  filter  or  adhering  to  the  receiver  is  dissolved  in  hydrochloric 
acid;  the  barium  is  precipitated  from  the  solution  as  sulphate, 
ignited  and  weighed;  and  from  the  weight  of  barium  sulphate 
is  calculated  the  equivalent  amount  of  carbon  dioxide  originally 
liberated. 

The  apparatus,  shown  in  the  figure,  consists  of  an  evolution 
flask  (F)  of  about  50  cm.3  capacity;  an  absorption  cylinder  (C) 
made  of  wide  glass  tubing;  a  toy  rubber 
balloon ;  rubber  stoppers,  and  glass  tubes 
for  connection.  The  tube  (A)  which  con- 
nects flask  and  cylinder  is  wide  enough 
(about  0.7  cm.  bore)  to  prevent  forma- 
tion of  bubbles  as  the  liquid  distils,  and 
for  further  protection  is  expanded  to  a 
bulb.  The  stopper  at  the  lower  end  of 
the  cylinder,  placed  vertically,  carries  a 
short  tube,  about  1.5  cm.  in  bore,  to 
which  the  rubber  balloon  is  securely 
bound.  The  stopper  at  the  upper  end 
of  the  cylinder  is  perforated  with  two 
holes,  through  one  of  which  passes  the 
tube  of  a  glass  stopcock,  while  through 
the  other  passes  a  long  tube  reaching  to 
the  interior  of  the  balloon,  and  provided 
with  a  Bunsen  valve  (preferably  of  Kreider's  pattern  f). 

In  using  this  apparatus,  a  saturated  solution  of  barium  hy- 
droxide (which  is  filtered  into  a  siphon  bottle,  and  preserved 
from  atmospheric  action  by  a  floating  layer  of  kerosene)  is 
introduced  by  pressure  upon  the  air  in  the  siphon  bottle  or  by 
suction  applied  to  the  stopcock  of  the  cylinder.  Such  a  solu- 
tion contains  about  5  per  cent  of  its  weight  of  the  hydroxide. 
It  is  best  to  use  of  it  in  every  case  an  amount  at  least  a  fourth 
in  excess  of  the  quantity  theoretically  required  to  absorb  the 

*  F.  A.  Gooch  and  I.  K.  Phelps,  Am.  Jour.  Sci.,  [3],  1,  101. 
t  See  page  7. 


230  METHODS  IN   CHEMICAL  ANALYSIS 

carbon  dioxide,  and  to  fill  the  cylinder  and  balloon  nearly  full 
of  liquid.  The  carbonate  is  weighed,  introduced  into  the  flask, 
and  washed  down  with  15  or  20  cubic  centimeters  of  boiled 
water,  which  is  protected  from  carbon  dioxide  of  the  breath  by  a 
balloon  attached  to  the  inlet  tube  within  the  wash  bottle,  so 
that  upon  blowing  into  the  inlet  tube  the  breath  distends  the 
balloon  and  thus  creates  the  necessary  pressure.  A  small  tube, 
holding  enough  hydrochloric  acid  to  effect  the  decomposition 
of  the  carbonate  to  be  analyzed,  is  placed  in  upright  position 
in  the  evolution  flask.  The  stopper  is  inserted  in  the  flask  and 
connections  are  made  as  shown  in  the  figure,  the  little  tube  con- 
taining the  acid  is  overturned  by  inclining  the  flask,  the  acid 
mixes  with  the  water,  and  effervescence  begins.  Heat  is  applied 
and  the  liquid  in  the  flask  is  boiled  until  that  in  the  cylinder  is 
heated  by  the  steam  nearly  to  the  boiling  point,  in  order  that  the 
precipitated  barium  carbonate  may  become  as  granular  as  pos- 
sible. The  carbon  dioxide  evolved  and  the  air  in  the  flask  are 
transferred  in, the  process  to  the  absorption  cylinder,  the  valve 
serving  to  prevent  the  back-flow  of  the  liquid  while  the  balloon 
expands  to  give  room  to  the  air  and  condensed  steam.  When  the 
boiling  is  done  the  flask  and  tube  are  disconnected  at  the  rubber 
joint,  the  cylinder  is  shaken  to  insure  the  absorption  of  the 
carbon  dioxide,  and  the  liquid  carrying  the  greater  part  of  the 
precipitate  is  transferred  through  the  stopcock  to  a  paper  filter 
moistened  with  water,  and  containing  about  5  cm.3  of  xylene. 
The  function  of  the  xylene,  which  was  found  to  be  preferable  to 
benzene,  kerosene  or  amyl  alcohol,  is  to  rise  to  the  surface  when 
the  aqueous  solution  is  added,  so  as  to  protect  the  barium  hy- 
droxide from  the  action  of  the  carbon  dioxide  of  the  air.  By 
manipulating  the  balloon  and  the  stopcock  (to  which  a  little 
funnel  may  be  attached  by  a  piece  of  rubber  tubing  for  con- 
venience in  introducing  wash  water)  the  cylinder  may  be  emptied 
and  washed  out  with  hot  boiled  water,  though,  of  course,  a  very 
considerable  portion  of  the  precipitate  remains  adhering  to  the 
walls  of  the  absorption  apparatus. 

The  filter  is  prepared  for  use  with  the  suction  pump,  but  in 
the  early  stages  of  filtration  and  washing  very  little  suction 
should  be  applied.  When  the  barium  hydroxide  has  been  nearly 
washed  out  of  the  precipitate,  the  xylene  is  dissolved  in  a  little 
hot  alcohol,  the  suction  is  applied  and  the  washing  is  completed 


CARBON 


231 


with  hot  water.  The  emulsion  of  xylene  and  water  found  in  the 
nitrate  is  readily  cleared  up  by  alcohol.  Finally,  the  barium 
carbonate  in  the  absorption  apparatus  and  upon  the  filter  is 
dissolved  in  hydrochloric  acid  and  precipitated  in  hot  solution 
by  sulphuric  acid,  the  resulting  barium  sulphate  is  filtered, 
washed,  and  ignited  upon  asbestos  in  a  perforated  crucible,  and 
from  its  weight  the  carbon  dioxide  which  originally  precipitated 
the  barium,  now  in  the  form  of  the  sulphate,  is  calculated.  The 
results  of  a  series  of  determinations  made  in  this  manner  are 
recorded  in  the  following  table. 

Carbon  Dioxide  by  Precipitation. 


CaCO3  taken, 
grm. 

BaSO«  found, 
grm. 

CO2  actually 
present. 

grm. 

CO2  calculated 
from  BaSO4. 

grm. 

Error  in  CO2. 
grm. 

O.O5OO 

o.  1180 

0.0220 

0.0222 

+O.OOO2 

0.0500 

0.1183 

0.0220 

0.0223 

+0.0003 

O.IOOO 

0.2329 

O.O44O 

0.0439 

—  o.oooi 

I.  1000 

0.2347 

o  .  0440 

O.O442 

+0.0002 

0.2000 

o  .  4660 

0.0880 

0.0878 

—  0.0002 

O.2OOO 

0.4653 

0.0880 

0.0876 

—0.0004 

0.5000 

I  .  1650 

O.22OO 

0.2196 

—0.0004 

0.5000 

1.1657 

O.22OO 

0.2197 

-0.0003 

I  .OOOO 

2.3323 

0.44OO 

o  .  4396 

—0.0004 

I.  0000 

2.3309 

0.4400 

o  •  4394 

—0.0006 

The  process  is  fairly  rapid  and  accurate. 

The  lodometric  Determination  of  Carbon  Dioxide. 

When  the  solution  of  a  metallic  hydroxide  is  acted  on  by 
iodine  at  a  temperature  high  enough  to  decompose  the  small 
amounts  of  hypoiodites  that  might  otherwise  be  present,  the  final 
action  results  in  the  formation  of  an  exactly  neutral  mixture  of 
iodate  and  iodide,  according  to  the  equation: 

6  ROH  +  3  I2  =  RI03  +  5  RI  +  3  H2O. 

Upon  the  assumption  that  in  the  case  of  barium  hydroxide  this 
reaction  is  regular  under  the  conditions  of  analysis  and  inde- 
pendent of  the  excess  of  iodine  which  remains  in  the  neutral 
mixture  unacted  upon,  and  that  iodine  may  be  estimated  by 
directly  titrating  with  arsenious  acid,  Phelps  *  has  elaborated  a 
differential  method  for  determining  carbon  dioxide.  In  this 
*  Am.  Jour.  Sci.,  [4],  ii,  70. 


232  METHODS  IN  CHEMICAL  ANALYSIS 

method  the  liberated  gas  is  run  into  a  measured  amount  of 
barium  hydroxide,  the  final  excess  of  which  is  estimated  by 
treating  with  iodine  in  the  presence  of  the  precipitated  barium 
carbonate. 

The  solution  of  barium  hydroxide  is  standardized  by  drawing 
80—90  cm.3  of  decinormal  iodine  into  a  glass  flask,  provided  with 
a  ground-glass  stopper  carrying  an  inlet  tube  reaching  nearly  to 
the  bottom  of  the  flask  and  an  outlet  tube  to  which  is  sealed 
a  Will  and  Varrentrapp  absorption  apparatus,  and  then  intro- 
ducing an  appropriate  amount  of  the  barium  hydroxide  solu- 
tion either  from  a  burette  or  from  a  stoppered  funnel  which  is 
weighed  before  and  after.  A  glass-stoppered  wash  bottle  answers 
for  a  standardizing  flask,  and,  with  the  glass  stopper  and  its 
attachments  replaced  by  a  rubber  stopper,  answers  the  pur- 
pose of  the  absorption  flask  described  later.  The  glass  stopper 
is  introduced,  the  inlet  being  closed  by  a  rubber  cap,  and  the 
absorption  apparatus  is  charged  with  a  solution  of  potassium 
iodide,  to  hinder  the  escape  of  iodine.  The  solution  is  heated  to 
boiling  to  break  up  traces  of  hypoiodite,  then  cooled  and  the 
excess  of  iodine  determined  by  decinormal  arsenite.  It  is  as- 
sumed that  the  iodine  lost  has  acted  on  the  barium  hydroxide 
according  to  the  equation : 

6  BaO2H2  +  6 12  =  Ba(IO3)2  +  5  BaI2  +  6  H2O. 

In  a  precisely  similar  manner  is  determined  the  quantity  of 
barium  hydroxide  remaining  after  a  measured  amount  of  it  has 
been  submitted  to  the  action  of  carbon  dioxide.  The  difference 
between  the  quantity  thus  formed  and  that  measured  out  is 
equivalent  to  the  carbon  dioxide  entering  into  action, 
carbon  Dioxide  A  convenient  apparatus  for  evolving  the  carbon 
in  Carbonates,  dioxide  from  carbonates  consists  of  a  wide-mouthed 
flask  of  about  75  cm.3  capacity,  furnished  with  a  doubly  perforated 
stopper  carrying  a  separating  funnel  for  the  introduction  of  acid 
into  the  flask,  and  a  tube  of  0.7  cm.  internal  diameter,  which  is 
expanded  to  a  small  bulb  just  above  the  stopper,  to  carry  off  the 
gas.  This  exit  tube  is  joined  by  a  rubber  connector  to  a  tube 
which  passes  through  the  rubber  stopper,  closing  the  absorption 
flask*  (the  glass-stoppered  wash  bottle  used  in  standardizing  the 
barium  hydroxide  solution),  and  which  ends  in  a  valve,  preferably 
*  See  left-hand  flask  in  Fig.  19,  page  237. 


CARBON  233 

of  the  Kreider  pattern.*  This  valve  is  inclosed  in  a  larger  tube 
reaching  nearly  to  the  bottom  of  the  absorption  flask.  Through 
a  second  hole  in  the  stopper  of  the  absorption  flask  passes  a  glass 
tube  closed  by  a  rubber  connector  and  screw  pinchcock. 

In  making  a  determination  of  carbon  dioxide  in  carbonates,  the 
weighed  substance  is  introduced  into  the  boiling  flask.  Barium 
hydroxide  solution,  in  amount  7  cm.3  to  10  cm.3  more  than 
actually  necessary  to  precipitate  the  carbon  dioxide,  is  drawn 
into  the  absorption  flask,  which  is  then  connected  with  the  boil- 
ing flask,  as  described  above.  The  stopcock  of  the  separating 
funnel  is  shut  off  and  the  flasks  evacuated  by  connecting  the  exit 
tube  of  the  absorption  flask  with  a  filter  flask  previously  pumped 
out  by  the  water  pump.  Phosphoric  acid  (chosen  as  a  nonvola- 
tile acid)  is  introduced  into  the  stoppered  funnel  with  about 
50  cm.3  of  water,  previously  purified  from  carbon  dioxide  by 
boiling  down  one-third,  and  kept  in  full,  stoppered  flasks  until 
used.  The  acid  is  allowed  to  enter  the  boiling  flask  and  the  car- 
bon dioxide  driven  over  completely  to  the  absorption  flask  by 
boiling  for  five  minutes  —  the  latter  being  shaken  frequently 
during  the  passage  of  the  gas  into  it  and  kept  cool  by  standing 
in  a  dish  of  water.  The  atmospheric  pressure  is  restored  by 
admitting  purified  air  through  the  funnel  of  the  boiling  flask. 
The  inlet  tube  of  the  absorption  flask  is  then  closed  by  a  rubber 
cap,  the  exit  tube  attached  to  potash  bulbs  and  the  flask  heated 
to  the  boiling  point  of  the  liquid  (to  granulate  the  precipitated 
carbonate)  and  then  cooled  in  a  stream  of  water.  The  exit  tube 
is  removed,  a  capillary  tube  long  enough  to  reach  below  the 
surface  of  the  liquid  introduced  and  decinormal  iodine  run  in 
until  the  liquid  is  yellow.  Then  the  glass  stopper  of  the  absorp- 
tion flask  is  introduced,  with  a  rubber  cap  on  the  inlet  tube  and 
potassium  iodide  solution  in  the  trap,  as  in  standardizing,  and 
the  emulsion  brought  to  a  boil.  Iodine  is  again  run  into  the  hot 
solution  through  the  inlet  tube  until  the  color  remains  distinctly 
red.  After  cooling,  the  excess  of  iodine  is  determined  by  stand- 
ard arsenious  acid. 

Results  obtained  by  treatment  of  calcite,  essentially  as  de- 
scribed, are  given  on  the  next  page,  correction  being  made  for 
the  deficiency  in  carbon  dioxide  (0.0014  grm.  for  each  gram  of 
the  mineral)  found  by  the  ignition  analysis. 

*  See  page  7. 


2,34 


METHODS  IN  CHEMICAL  ANALYSIS 


Carbon  Dioxide  by  Absorption  in  Barium  Hydroxide  and  lodometric  Determi- 
nation of  the  Excess. 


Calcite 
taken, 
grm. 

BaOjH2  taken, 
grm. 

BaO2H2  found, 
grm. 

CO2  found, 
grm. 

Error  on  CO2. 
grm. 

Error  corrected, 
grm. 

0.0501 

0.2484 

0.1604 

0.0227 

+0.0006 

+0.0007 

o  .  0500 

0.2381 

0.1508 

0.0224 

+0.0004 

+0.0005 

O.IO22 

0.3416 

0.1675 

o  0447 

—0.0003 

—o.oooi 

O.IO26 

0.3105 

0.0351 

0.0450 

—  o.oooi 

o.oooo 

O.2O32 

0.6181 

0.2692 

0.0896 

+O.OOO2 

+0.0004 

o  .  2049 

0.5761 

0.2223 

o  .  0908 

+O.OOO6 

+0.0008 

0.5088 

1.1301 

o  .  2606 

.     0.2232 

—  O.OOO7 

o.oooo 

0-50I5 

i  .  0804 

0.2245 

0.2197 

—  O.OOIO 

—0.0003 

1.0032 

2.0125 

0.3004 

0-4394 

—  O.OO2O 

—0.0006 

I.OO64 

2.0702 

0.3538 

0.4405 

—  0.0023 

—0.0009 

The  Combustion  of  Organic  Substances  in  the  Wet  Way. 

Phelps  *  has  studied  the  combustion  of  substances  in  the  wet 
way,  and,  with  apparatus  like  that  previously  described  for  the 
determination  of  the  carbon  dioxide  of  carbonates,  has  shown 
that  the  carbon  of  many  organic  compounds  may  be  determined 
by  oxidation  and  estimation  of  carbon  dioxide, 
carbon  content  Certain  organic  material  may  be  oxidized  by  po- 
bythePerman-  tassium  permanganate  in  acid  or  alkaline  solution, 
ganate  Process.  an(^  ^  car]I>on  COntent  of  the  organic  matter  may 
be  estimated  by  absorbing  in  a  measured  amount  of  standardized 
barium  hydroxide  the  carbon  dioxide  produced,  acting  upon  the 
residual  hydroxide  with  standard  iodine  in  excess,  and  deter- 
mining by  decinormal  arsenite  the  excess  of  iodine.  According 
to  this  process,  the  organic  material  is  weighed  out  and  intro- 
duced into  an  evolution  flask  with  10  cm.3  to  15  cm.3  of  pure 
water  (freed  from  carbon  dioxide  by  boiling  down  one-third  and 
kept  in  stoppered  bottles),  and  the  evolution  flask  is  connected 
with  an  absorption  flask  charged  with  a  volume  of  standardized 
barium  hydroxide  3  cm.3  to  5  cm.3  in  excess  of  the  amount  re- 
quired to  precipitate  the  carbon  dioxide  to  be  determined.! 

The  whole  system  is  then  evacuated  with  the  water  pump  to 
a  pressure  of  200  mm.  to  225  mm.  and  the  boiling  flask  warmed. 

*  I.  K.  Phelps,  Am.  Jour.  Sci.,  [4],  iv,  372. 

t  See,  for  the  general  arrangement,  the  more  elaborate  apparatus  of  Fig.  19 
.page  237. 


CARBON 


235 


An  excess  of  potassium  permanganate  solution  (prepared  by  dis- 
solving in  water,  acidulating  with  sulphuric  acid  and  boiling  until 
free  from  carbon  dioxide)  is  then  run  in  through  the  funnel  tube 
and  the  mixture  warmed  again.  The  carbon  dioxide  is  set  free 
by  sulphuric  acid  [i  13],  either  at  once  or  after  introducing 
an  excess  of  pure  sodium  hydroxide  and  boiling,  and  driven 
completely  to  the  absorption  flask  by  boiling  for  five  minutes. 
During  the  passage  of  the  gas  the  absorption  flask  is  shaken 
frequently  and  kept  cool  by  standing  in  a  dish  of  water  and  by 
pouring  cold  water  over  it  from  time  to  time.  During  the  boil- 
ing, the  vacuum  in  the  flasks  may  be  tested  by  opening  momen- 
tarily the  stopcock  of  the  funnel  tube  and  noting  the  direction 
of  the  flow  of  water  contained  in  the  funnel.  After  the  boiling, 
the  atmospheric  pressure  is  restored  by  allowing  air,  purified 
from  carbon  dioxide  by  passage  through  potash  bulbs,  to  enter 
through  the  funnel  tube  of  the  boiling  flask.  The  flasks  are 
disconnected,  the  stopper  of  the  absorption  flask  with  its  attach- 
ments is  removed  and  carefully  washed  free  from  barium  hydrox- 
ide, and  a  second  stopper,  provided  with  a  separating  funnel 
and  a  Will  and  Varrentrapp  absorption  apparatus  containing 
water  to  serve  as  a  trap,  is  inserted  into  the  mouth  of  the  absorp- 
tion flask.  The  contents  of  the  flask  are  brought  to  the  boiling 
point,  decinormal  iodine  solution  is  run  in  through  the  funnel 
tube  in  sufficient  quantity  to  destroy  the  larger  part  of  the  excess 
of  barium  hydroxide,  the  mixture  again  heated  to  boiling,  and 
iodine  run  in  again  to  permanent  red  coloration.  After  cooling, 
the  excess  of  iodine  is  determined  by  titration  with  decinormal 
arsenite.  The  difference  between  the  barium  hydroxide  taken 
and  that  equivalent  to  the  iodine  which  has  disappeared  is  the 

Oxidation  with  Permanganate  in  Acid  Solution. 


Ammonium 
oxalate  taken, 

grm. 

BaO,H,  taken, 
grm. 

BaO2H, 

remaining. 

grm. 

CO2  found, 
grin. 

CO, 

calculated. 

grm. 

Error  on  COj. 

grrn. 

0.2522 

0.7267 

o  .  i  i  70 

0.1565 

0.1561 

+o  .  0004 

0.2542 

0.7267 

O.III3 

0.1579 

0.1574 

+o  .  0005 

0.5020 

i  -4535 

0.2417 

0.3110 

0.3108 

+O.OOO2 

0.5058 

1-3954 

0.1753 

0.3131 

0.3131 

0.0000 

1.0033 

2.6163 

0.1955 

0.6213 

0.6211 

+0.0002 

1.0003 

2.5951 

0.1836 

0.6189 

0.6192 

-0.0003 

I.OOIO 

2.6163 

0.2037 

0.6192 

0.6197 

—  0.0005 

236 


METHODS  IN  CHEMICAL  ANALYSIS 


Oxidation  with  Permanganate  in  Alkaline  Solution. 


Barium 
formate  taken. 

BaO2H2  taken. 

BaO2H2 

remaining. 

CO2  found. 

CO, 

calculated. 

Error  on  COj. 

grm. 

grm 

grm 

grm. 

grm. 

grm. 

0.5001 

0.9302 

0.1745 

0.1939 

0.1935 

+O  .  0004 

0.5033 

0.9012 

o.  1402 

0.1953 

0.1947 

+0.0006 

I  .  OOO2 

1.6861 

0.1793 

0.3867 

0.3870 

—  0.0003 

I  .  0059 

1.6279 

0.1093 

0.3897 

0.3892 

+O  .  0005 

1-3750 

2.2529 

o.  1820 

0.5315 

0.5320 

—  0.0005 

1.5028 

2.4419 

0.1754 

0.5816 

0.5814 

+  0.0002 

Tartar  emetic 
taken.* 

grm. 

BaO2H2  taken, 
grm. 

BaO,H2 
remaining. 

grm. 

CO2  found, 
grm. 

C02 

calculated. 

grm. 

Error  on  COj. 
grm. 

0.5051 

i  .  2450 

0.1709 

0.2756 

0.2751 

+0.0005 

0.5030 

i  .2226 

O.IS36 

0.2743 

0.2739 

+0  .  0004 

0.7509 

1-7355 

o.  1401 

0.4094 

0.4091 

+0.0003 

0.7541 

i  •  7430 

o.  1410 

0.4111 

0.4107 

+0.0004 

i  .0018 

2.3456 

0.2187 

0.5458 

0.5456 

+  O.OOO2 

1.0005 

2.2435 

o.  1196 

0.5451 

0.5450 

+O.OOOI 

*  Dried  at  100°. 

measure  of  the  carbon  dioxide.  The  figures  given  show  the 
result  of  the  permanganate  oxidation  when  applied  to  ammo- 
nium oxalate  in  acid  solution,  and  to  barium  formate  and  tartar 
emetic  in  alkaline  solution. 

Obviously,  certain  organic  substances  oxidizable  by  perman- 
ganate may  be  analyzed  by  the  process  outlined  above,  but, 
as  noted  by  Wanklyn  and  Cooper*  and  by  others,  potassium 
permanganate,  whether  in  acid  or  alkaline  solution,  fails  to  oxi- 
dize completely  certain  organic  substances  (acetates  and  carbo- 
hydrates, for  example)  even  at  the  boiling  temperature. 
Carbon  Content  It  is  well  known,  however,  that  a  mixture  of  con- 
whhCnromix:  centrated  sulphuric  and  chromic  acids  has  a  much 
Acid.  wider  field  of  action  in  oxidizing  organic  compounds, 

and  Phelpsf  has  studied  the  application  of  this  method  also.  The 
apparatus  recommended  for  the  process,  and  shown  in  the  ac- 
companying figure,  may  be  described  as  follows: 

A  thick-walled  liter  flask  with  round  bottom,  serving  as  an 
oxidizing  chamber,  is  closed  by  a  rubber  stopper  with  two  per- 

*  Phil.  Mag.,  [5],  vii,  138. 
t  Loc.  cit. 


CARBON 


237 


1-ig.    n 


forations.  Through  one  of  these  passes  the  tube  of  a  separating 
funnel  reaching  nearly  to  the  bottom  of  the  flask  and  drawn  out 
at  the  lower  end .  A  disk  of  platinum 
foil,  through  which  the  tube  of  the 
funnel  passes,  is  hung  in  the  neck  of 
the  flask,  nearly  closing  it,  and  held 
in  place  by  an  attached  platinum 
wire,  the  end  of  which  is  squeezed 
under  the  rubber  stopper.  Through 
the  second  hole  of  the  stopper  passes 
an  exit  tube  0.7  cm.3  in  bore.  This 
tube,  expanded  just  above  the  stop- 
per to  a  small  bulb  which  serves  to 
prevent  mechanical  loss  of  the  solid 
contents  of  the  flask  during  the  boil- 
ing, is  joined  by  means  of  a  rubber 
connector  (provided  with  a  screw 
pinchcock)  to  the  inlet  tube  of  the  absorption  flask,  which  is 
an  ordinary  500  cm.3  round -bottomed  flask.  This  flask  is  also 
closed  by  a  rubber  stopper  with  two  perforations,  through  one 
of  which  passes  the  inlet  tube  and  through  the  other  the  exit 
tube,  which  is  also  enlarged  to  a  small  bulb  just  above  the  stop- 
per and  is  closed  by  a  rubber  connector  and  screw  pinchcock. 
The  ground-glass  stopper  of  the  funnel  tube  is  carefully  cleaned 
and  lubricated  with  a  thick  solution  of  metaphosphoric  acid. 

A  partial  vacuum  is  easily  obtained  by  boiling  water  in  the 
evolution  flask  and  the  barium  hydroxide  solution  in  the  absorp- 
tion flask  at  the  same  time,  closing  the  flasks,  and  cooling — both 
flasks  being  connected  and  ready  for  a  determination. 

In  making  a  determination,  the  organic  substance  is  weighed 
out  in  a  counterbalanced  bulb,  so  thin  that  it  may  be  easily 
broken  later  and  made  with  a  wide  mouth  for  convenience  in 
introducing  the  solid  substance.  After  the  substance  is  weighed, 
the  mouth  of  the  bulb  is  sealed  by  heating  in  a  small  blowpipe 
flame  and  the  tube  introduced  into  the  evolution  flask,  together 
with  an  amount  of  pure  potassium  dichromate,  known  to  be  in 
excess  of  that  required  to  oxidize  the  organic  substance.  The 
flasks  are  connected,  as  already  described,  with  an  appropriate 
amount  of  barium  hydroxide  solution  in  the  absorption  flask  and 
10  cm.3  of  pure  water  in  the  evolution  flask,  and  the  vacuum 


238 


METHODS  IN  CHEMICAL  ANALYSIS 


obtained  (as  described  above)  by  boiling  the  contents  of  both 
flasks  until  the  water  in  the  evolution  flask  has  decreased  to  2  cm.3 
or  3  cm.3,  and  cooling.  The  tube  containing  the  organic  sub- 
stance is  then  broken  by  shaking  the  flask,  and  20  cm.3  of  concen- 
trated sulphuric  acid,  previously  purified  from  organic  material 
by  heating  to  the  fuming  point  with  a  few  crystals  of  potas- 
sium dichromate,  are  run  in  through  the  funnel  tube.  While 
still  hot,  the  acid  is  shaken  in  the  flask  violently,  the  platinum 
foil  hung  in  the  neck  serving  to  protect  the  rubber  stopper.  The 
flask  is  warmed  to  approximately  105°,  the  highest  temperature 
to  which,  as  shown  by  Cross  and  Bevan,*  such  a  mixture  of 
chromic  and  sulphuric  acids  may  be  safely  heated  without  the 
disengagement  of  oxygen  gas.  Water  is  then  run  in  until  the 
crystals  of  chromic  anhydride  have  disappeared  and  the  danger 
of  the  evolution  of  oxygen  is  past.  The  solution  is  heated  to  its 
boiling  point,  with  care  to  keep  the  outward  pressure  less  than 
the  inward  pressure  —  the  relation  being  easily  observed  by 
opening  momentarily  the  stopcock  of  the  funnel  tube  and  noting 
the  direction  of  the  flow  of  water  contained  in  the  funnel.  The 
flask  is  shaken  and  heated  alternately  for  five  minutes  to  bring 
about  the  oxidation  of  small  amounts  of  carbon  monoxide, 
originally  produced.  Then  more  water  (60  cm.3  to  70  cm.3)  is 
introduced  through  the  funnel  and  the  stopcock  between  the 

Oxidation  with  Chromic  Acid. 


Substance 
taken. 

Ba02H2 
taken. 

Ba02H2 
found. 

CO, 

found. 

C02 

calculated. 

Error  on 
C02. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

Analysis  of  ammonium  oxalate. 


0.5009 

1-3534 

o.  1469 

0.3097 

0.3101 

—  0.0004 

0.5006 

1.3400 

0.1308 

0.3103 

0.3099 

+o  .  0004 

0.5005 

1.3400 

0-1343 

0.3094 

0.3098 

—  0.0004 

I  .0002 

2  .  5460 

0.1347 

0.6188 

0.6192 

—0.0004 

I  .0010 

2.5192 

'  o.  1094 

0.6185 

0.6197 

—  0.0012 

Analysis  of  cane  sugar. 


O.  2OOI 

1.3926 

0.1905 

0.3085 

O^OSS 

—  0.0003 

0.2000 

1.3926 

0.1936 

0.3077 

0.3086 

—  0.0009 

O.2OOI 

1.3926 

0.1857 

0.3097 

0.3088 

+  0.0009 

O.2OI4 

1.3400 

0.1279 

0.3III 

0.3108 

+0.0003 

*  Jour.  Chem.  Soc.,  liii,  889. 


CARBON  239 

boiling  and  absorption  flasks  is  opened  to  admit  the  carbon 
dioxide  to  the  latter,  which  is  kept  cool  and  shaken  as  before. 
The  contents  of  the  evolution  flask  are  then  heated  to  boiling 
and  a  slow  current  of  air,  freed  from  carbon  dioxide  by  passage 
through  potash  bulbs,  is  allowed  .to  enter  through  the  funnel  tube 
to  keep  the  liquid  from  undue  bumping.  The  boiling  is  con- 
tinued for  fifteen  minutes,  after  which  the  excess  of  barium  hy- 
droxide is  determined  iodometrically  and  the  carbon  dioxide 
calculated. 

The  results  of  experiments  with  ammonium  oxalate  and  with 
cane  sugar,  one  of  the  more  difficultly  oxidizable  substances, 
are  given  in  the  preceding  table.  The  results  are  evidently 
very  satisfactory. 

carbon  Dioxide  Taking  advantage  of  the  fact  that  at  105°  the 
Evolved  and  mixture  of  chromic  and  sulphuric  acids  does  not 
oxygen  Used.  evoive  OXygen,  Phelps*  has  been  able  to  determine 
the  oxygen  used  in  the  combustion  of  certain  organic  substances. 
The  knowledge  of  this  amount  of  oxygen  and  of  that  contained 
in  the  products  of  oxidation  gives,  by  difference,  the  oxygen 
content  of  the  original  substance. 

In  this  operation  a  known  weight  of  the  substance  is  first 
treated  with  a  known  weight  of  pure  potassium  dichromate  and 
20  cm.3  of  concentrated  and  purified  sulphuric  acid,  according  to 
the  process  described  above  for  the  determination  of  carbon  as 
carbon  dioxide.  The  residual  liquid,  which,  after  dilution  and 
boiling  for  the  removal  and  estimation  of  the  carbon  dioxide, 
should  have  a  volume  of  60  cm.3  to  80  cm.3,  is  washed  into  the 
Voit  flask  of  the  distillation  and  absorption  apparatus  figured 
and  previously  described.!  The  absorption  chamber  is  charged 
with  a  solution  of  arsenious  oxide  in  sodium  hydroxide,  the  for- 
mer being  in  slight  excess  of  the  amount  required  to  take  up  the 
chlorine  to  be  evolved  by  the  chromate  and  the  latter  in  quantity 
more  than  sufficient  to  neutralize  the  acid  which  may  be  volatil- 
ized to  the  receiver  in  the  later  operation.  The  apparatus  is 
connected,  hydrochloric  acid  (35  cm.3  of  the  strongest  acid) 
introduced  through  the  stoppered  funnel,  a  slow  current  of 
carbon  dioxide  started  through  the  system,  and  the  liquid  in 
the  flask  slowly  boiled  down,  for  a  period  of  five  or  six  hours, 

*  Am.  Jour.  Sci.,  [4],  ii,  379. 
t  See  Fig.  3,  page  4. 


240 


METHODS  IN  CHEMICAL  ANALYSIS 


until  the  volume  is  30  cm.3  to  40  cm.3.  After  cooling  and  dis- 
connecting the  apparatus,  the  solution  in  the  receiver  is  made 
acid  with  sulphuric  acid  and  then  alkaline  with  acid  potassium 
carbonate.  The  residual  arsenite  is  determined  by  titration  with 
decinormal  iodine. 

Determination  of  Carbon  Dioxide  Evolved  and  Oxygen  Used. 


Oxygen 

Substance 
taken. 

C02 
found. 

Error  on 
C02. 

K2Cr207 
taken. 

As2O3 
taken. 

As203 
found. 

Oxygen 
used. 

required 
by 

Error  on 
oxygen. 

theory. 

grm.. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

Analysis  of  ammonium  oxalate. 


I.OI22 
I  .OOI9 

0.6265 
0.6212 

—  O.OOOI 
+0.0010 

2.0009 

2.O002 

1.3002 
I-3SI7 

o  .  oooo 
o  .  0440 

0.1160 
0.1147 

0.1139 
0.1128 

+O.OO2I 
+O.OOI9 

Analysis  of  phthalic  acid. 


O.IOO2 
0.1093 

0.2138 
0.2324 

+O.OOI4 
+O.OO07 

2.OOI2 
2.  OOOO 

1.2004 
I.I03I 

0.0814 
0.0634 

0.1456 
0.1582 

0.1448 
0.1580 

+0.0008 
+O.OOO2 

Analysis  of  cane  sugar. 


0.2025 
0.4012 

0.3117 

0.6166 

—0.0008 
—0.0024 

3  .  oooo 
5.0000 

i  .  7002 
2.3022 

0.0796 
0.0366 

0.2275 
0.4495 

0.2273 
0.4502 

+0  .  0002 
—  0.0007 

Analysis  of  paper. 


0.3034 
0.4523 

0.4932 
0.7334 

—  O.OOIO 

-0.0033 

3.5015 

5-0035 

I  .4017 
I  .  8000 

0.0879 
0.0710 

0.3539 
0.5368 

0.3598 
0.5358 

—0.0005 

+O.OOIO 

Analysis  of  tartar  emetic. 


0.5057 
1.0099 

0.2671 

0.5321 

—0.0009 
—0.0030 

2.5018 

3-5003 

I  .7000 

1.7520 

0.0766 
0.0198 

0.1459 

0.2911 

0.1462 
0.2919 

-0.0003 

—0.0008 

Analysis  of  barium  formate. 


1.0079 
1.5014 

0.3906 
0.5814 

+0.0006 
+0.0005 

3.0026 
3.0010 

2  .  2002 
I  .  8080 

o  .  0496 
o  .  0890 

0.1423 

0.2118 

0.1422 
0.2118 

+O.OOOI 

o  .  oooo 

When  the  boiling  begins  the  chromate  is  gradually  reduced 
with  evolution  of  chlorine,  but  if  the  evaporation  of  water  is 
pushed  too  rapidly  the  sulphuric  acid  may  reach  a  degree  of  con- 


SILICON  241 

centration  such  that  oxygen  is  evolved  instead  of  chlorine  in  the 
process  of  reduction.  Sometimes,  during  the  reduction,  chloro- 
chromic  anhydride  is  visibly  volatilized  to  the  receiver,  but  inas- 
much as  it  is  there  reduced  and  registered  by  the  arsenite  this 
transfer  is  of  no  moment. 

The  results  of  experiments  in  which  the  carbon  of  various 
organic  substances  was  first  determined  by  oxidation  according 
to  the  process  previously  described,*  the  amount  of  potassium 
dichromate  employed  being  accurately  known,  and  the  residual 
dichromate  found  by  reduction  with  hydrochloric  acid  by  the 
process  just  outlined,  are  given  in  the  tabular  statement. 

From  these  results,  it  will  be  seen  that  the  process  works  with 
accuracy  for  a  great  variety  of  organic  substances.  It  was  found 
impossible,  however,  to  determine  the  elements  in  compounds 
which,  like  ether  and  naphthalene,  are  at  the  same  time  volatile 
and  hard  to  oxidize  completely. 

SILICON. 
The  Detection  of  Silicon  in  Silicates  and  Fluo silicates. 

The  formation  of  silicon  fluoride  by  the  action  of  hydrofluoric 
acid  or  a  fluoride  and  sulphuric  acid  upon  a  silicate  is  often 
applied  to  the  detection  of  silica,  the  silicon  fluoride  giving  with 
water  a  white  precipitate  of  silicic  acid.  The  usual  proce- 
dure in  making  this  test  is  to  allow  the  gaseous  silicon  fluoride  to 
come  in  contact  with  a  moistened  glass  rod,  but  the  condensa- 
tion of  steam  or  sulphuric  acid  on  the  rod  often  makes  the  results 
uncertain.  Browningf  has  found  that  when  moistened  black 
paper  is  brought  in  contact  with  the  fumes  of  the  gaseous  fluoride 
the  deposit  of  silica  is  very  easily  detected. 
According  to  the  procedure  recommended,  the 
reaction  is  made  to  take  place  in  a  small  lead 
cup  about  one  centimeter  in  diameter  and  depth, 
made  by  running  the  melted  metal  into  a  mold, 
covered  by  a  flat  piece  of  lead  with  a  small 
hole  in  the  center,  as  shown  in  Fig.  20.  Into  Flg-  20> 

this  cup  is  put  the  fluosilicate,  or  the  silicate  with  a  small  amount 
of  finely  powdered  calcium  fluoride,  generally  about  o.i  grm.,  and 
a  few  drops  of  concentrated  sulphuric  acid  are  added.     Upon  the 
*  See  page  236. 
t  Philip  E.  Browning,  Am.  Jour.  Sci.,  [4],  xxxii,  249. 


242 


METHODS  IN   CHEMICAL  ANALYSIS 


upper  side  of  the  cover  a  piece  of  moistened  black  filter  paper  is 
placed  and  upon  this  a  small  moistened  pad  of  ordinary  filter 
paper  to  keep  the  black  paper  moist  during  heating  upon  a 
steam  bath.  At  the  conclusion  of  the  heating  of  about  ten 
minutes,  a  white  deposit  is  found  on  the  under  side  of  the  black 
paper,  over  the  opening  in  the  cover,  if  silica  is  present  in  ap- 
preciable amount. 

Tests  of  this  procedure  are  given  in  the  table. 

Silicate  Tests. 


Material  tested, 
gnu  . 

Approximate 
per  cent  of  SiO2. 

CaF2  present, 
grin. 

Result. 

o  1000  SiOa  

IOO 

O.IOOO 

Nothing. 

o  1000  SiO2 

IOO 

O   IOOO 

Very  good. 

o  oioo  SiO2  

IOO 

O.  IOOO 

Very  good. 

o  0050  SiO2 

IOO 

O.  IOOO 

Very  good. 

o  ooio  SiO2                              

IOO 

O.IOOO 

Trace. 

o.oioo  Kaolinite  
0.0050  Kaolinite  
o  ooio  Kaolinite 

46 
46 

46 

O.  IOOO 
O.  IOOO 
O.  IOOO 

Very  good. 
Very  good. 
Trace. 

o  oioo  Gadolinite 

24 

O.  IOOO 

Very  good. 

o  0050  Gadolinite 

24 

O.  IOOO 

Trace. 

o.oioo  Lepidolite  

50 

O.  2OOO 

Good. 

Fluosilicate  Tests. 


o  0050  Na2SiFe        

Very  good. 

o  ooio  Na2SiFg 

Very  good. 

TITANIUM. 

The  Determination  of  Titanic  Acid  by  Reduction  and  Titration 
with  Potassium  Permanganate. 

The  estimation  of  titanic  acid  by  reduction  with  zinc  and 
direct  titration  leads  to  low  results  even  when  precautions  are 
taken  to  avoid  atmospheric  oxidation  during  titration.*  New- 
tonf  has  shown,  however,  that  it  is  possible  to  determine  titanic 
acid  successfully  by  reducing  with  zinc  in  an  atmosphere  of 
hydrogen,  adding  an  excess  of  ferric  sulphate,  and  titrating  the 
resulting  and  equivalent  ferrous  salt  with  potassium  perman- 

*  Cf.  Pisani,  Compt.  rend.,  lix,  298;  Marignac,  Zeit.  anal.  Chem.,  vii,  112; 
Wells  and  Mitchell,  Jour.  Am.  Chem.  Soc.,  xvii,  878. 
t  H.  D.  Newton,  Am.  Jour.  Sci.,  [4],  xxv,  130. 


TITANIUM 


243 


ganate;    the  reaction  between   the  salts  of   titanium   and  iron 
taking  place  according  to  the  equation: 

Ti2(SO4)3  +  Fe2(SO4)3  =  2  Ti(SO4)2  +  2  FeSO4. 

According  to  this  procedure,  titanic  acid  dissolved  in  con- 
centrated sulphuric  acid  is  introduced  into  a  ioo-cm.3  flask  and 
the  solution  is  diluted  with  water  until  it  contains  10  per  cent 
of  sulphuric  acid,  this  strength  being  sufficient  to  hold  the  titanic 
acid  in  solution,  while  insufficient  to  reoxidize  reduced  titanium 
oxide.  Zinc  is  added  in  suitable  amount,  and  a  rubber  stopper 
carrying  a  delivery  tube  and  a  small  separating  funnel  is  inserted 
in  the  neck  of  the  flask.  After  the  air  has  been  driven  from 
the  flask  by  the  hydrogen  evolved,  the  delivery  tube  is  dipped 
under  water  and  the  stopcock  of  the  funnel  is  closed.  Gentle 
heat  is  applied  until  all  the  zinc  is  dissolved,  and  the  solution 
is  cooled.  An  excess  of  ferric  sulphate  is  passed  into  the  flask 
through  the  separating  funnel  and  followed  at  once  by  cold, 
freshly  distilled  water  until  the  flask  is  filled  to  the  neck.  The 
contents  of  the  flask  are  poured  into  more  cold  distilled  water 
and  the  ferrous  salt  produced  by  action  of  the  ferric  salt  upon  the 
reduced  titanium  oxide  is  titrated  by  n/io  permanganate. 

Test  results  are  given  in  the  table. 

Permanganate  Titration  of  Titanic  Acid  Reduced  by  Zinc. 


KMn04. 
cm.3 

TiO2  taken, 
grm. 

TiO2  found, 
grm. 

Error, 
grm. 

6.50 

0.0520 

0.0523 

+o  .  0003 

6.52 

0.0520 

0.0524 

+O.0004 

6-45 

0.0520 

0.0519 

—  o.oooi 

6.50 

0.0520 

0.0523 

+o  .  0003 

6.48 

0.0520 

0.0521 

+O.OOOI 

6.42 

0.0520 

0.0518 

—  O.OOO2 

6.50 

0.0520 

0.0523 

+0.0003 

19-95 

0.1596 

0.1599 

+0.0003 

19-93 

0.1596 

0.1598 

+O.OOO2 

19-95 

o.  1596 

0-1599 

+0.0003 

19.90 

o.  1596 

0.1595 

—o.oooi 

19-95 

o.  1596 

0-1599 

+0.0003 

19-95 

o.  1596 

0.1599 

+0.0003 

19.90 

0.1596 

0-1595 

—  O.OOOI 

19.85 

0.1596 

0.1591 

-0.0005 

19-95 

0.1596 

0.1599 

+0.0003 

19.88 

0.1596 

0.1594 

—  O.OOO2 

19-95 

0.1596 

0.1599 

+0.0003  • 

19-95 

0.1596 

0.1599 

+0.0003 

244  METHODS  IN  CHEMICAL  ANALYSIS 

ZIRCONIUM. 

The  Separation  of  Zirconium  from  Iron  by  Volatilization  of  the 
Latter  in  Hydrogen   Chloride. 

It  has  been  shown  by  Havens  and  Way*  that  zirconium  oxide 
may  be  separated  from  ferric  oxide  by  the  volatilization  of  ferric 
chloride  in  an  atmosphere  of  hydrogen  chloride,  containing  a 
little  chlorine,  at  a  temperature  of  200°  to  300°.  f 


CERIUM. 

The  Separation  of  Cerium  from  Other  Cerium  Earths  by  the  Action 
of  Bromine  upon  the  Mixed  Hydroxides  in  Presence 
of  an  Alkali  Hydroxide. 

One  of  the  best  known  processes  for  the  separation  of  cerium 
from  lanthanum  and  didymium  is  that  of  Mosander.J  This 
process  consists  in  passing  chlorine  gas  jnto  a  mixture  of  the 
hydroxides  suspended  in  a  distinct  excess  of  a  fixed  alkali  hy- 
droxide, until  the  solution  is  saturated  and  the  reaction  of  the 
liquid  is  no  longer  alkaline  to  litmus.  Under  these  conditions 
nearly  all  the  cerium  remains  undissolved  as  the  eerie  hydroxide, 
while  the  other  cerium  earths  go  largely  into  solution.  In  treat- 
ing mixed  material,  the  residue  of  eerie  hydroxide  generally 
retains  some  of  the  cerium  earths,  so  that  the  treatment  with 
chlorine  must  be  repeated.  Two  disadvantages  associated  with 
this  method,  therefore,  are  the  preparation  and  use  of  chlorine 
gas,  and  the  solvent  action  of  the  hydrochloric  acid  formed  in 
the  reaction  upon  the  eerie  hydroxide 

2  Ce(OH)3  +  C12  =  2  CeO2  +  2  HC1  +  2  H2O. 

Browning  and  Roberts  §  have  shown  that  by  substituting  bro- 
mine for  chlorine  in  the  Mosander  process  about  50  per  cent  of 
the  other  cerium  earths  can  be  separated  from  eerie  hydroxide 
in  one  treatment,  and  that  after  three  treatments  practically  all 
the  other  cerium  earths  are  removed  without  any  solvent  action 

*  Franke  Stuart  Havens  and  Arthur  Fitch  Way,  Am.  Jour.  Sci.,  [4],  viii,  217. 

t  See  page  508. 

{  Jour,  prakt.  Chem.,  xxx,  267. 

§  Philip  E.  Browning  and  Edwin  J.  Roberts,  Am.  Jour.  Sci,  [4],  xxix,  45. 


CERIUM 


245 


upon  the  eerie  hydroxide.  The  advantages  of  the  method  are, 
the  convenience  in  the  use  of  the  bromine,  and  the  apparent  lack 
of  tendency  of  the  hydrobromic  acid  to  dissolve  the  eerie  hy- 
droxide. 

The  procedure  is  as  follows :  The  mixed  hydroxides  are  precip- 
itated with  a  slight  excess  of  sodium  hydroxide  or  potassium  hy- 
droxide, and,  suspended  in  the  alkaline  solution,  are  treated  with 
liquid  bromine  or  bromine  water  in  distinct  excess,  and  the 
mixture  is  placed  upon  a  steam  bath  until  the  greater  part  of  the 
free  bromine  is  expelled.  The  residue  is  then  filtered  off,  washed, 
and  treated  as  before.  This  process  is  repeated. 

In  the  test  experiments  upon  a  mixture  composed  of  about 
50  per  cent  of  cerium  oxide  and  50  per  cent  of  cerium  earth  oxides 
other  than  cerium  oxide,  the  filtrate  after  each  treatment  was 
found  to  contain  the  amounts  of  cerium  earth  oxides,  free  from 
cerium,  indicated  in  the  table.  The  residue  from  the  last  treat- 
ment on  being  dissolved  in  acid  showed  only  faint  didymium 
bands.  In  another  experiment  a  larger  amount  of  material, 
10  grm.,  was  subjected  to  a  fourth  and  fifth  treatment  with 
bromine,  the  fourth  treatment  yielding  a  small  fraction  of  a 
gram  of  oxides  other  than  cerium  oxide,  and  the  fifth  only  a 
few  milligrams.  In  both  cases  these  oxides  were  free  from 
cerium.  The  oxides  from  the  first  filtrates  were  much  lighter 
in  color  than  those  obtained  from  the  last,  which,  of  course, 
indicates  that  the  lanthanum  is  dissolved  by  the  action  of  the 
bromine  more  readily  than  the  didymium.  The  quantitative 
results  follow  in  the  table. 

Separations  by  Bromine  and  Alkali  Hydroxide. 


Mixed  oxides 
taken. 

Oxides  found  in 
first  filtrate. 

Oxides  found  in 
second  filtrate. 

Oxides  found  in 
third  filtrate. 

Total  oxides 
found. 

grm. 

grm. 

grm. 

grm. 

grm. 

I  .OOOO 

0.3310 

0.0720 

0.0190 

0.4420 

I  .OOOO 
I.  0000 
I  .OOOO 
IO.OOOO 
IO.OOOO 

0.2900 
0.2250 
0.2750 
3.1360 
3-4590 

O.  IOIO 

o  .  i  290 

0.0860 
1.0050 
0.5240 

o  .  0420 
o  .  0640 
o  .  0740 

0.5930 
0.8560 

0.4330 
0.4180 
0.4350 
4-7340 
4.8390 

The  action  of  iodine  is  similar  to  that  of  chlorine  and  bromine, 
but  it  is  too  incomplete  to  be  of  practical  value. 


246  METHODS  IN  CHEMICAL  ANALYSIS 

The  lodometric  Estimation  of  Cerium. 

Anhydrous  cerium  dioxide,  prepared  by  the  ignition  of  the 
oxalate  or  hydroxide,  is  very  slowly  acted  on  by  acids,  especially 
when  pure.  For  this  reason  the  method  which  Bunsen  de- 
scribed has  remained  the  only  one  adapted  to  the  satisfactory 
volumetric  estimation  of  the  ignited  dioxide.*  According  to 
this  method,  the  substance  to  be  determined  is  weighed  out  in 
a  glass  flask  of  10  to  15  cubic  centimeters'  capacity,  a  few  crys- 
tals of  potassium  iodide  are  added,  and  the  neck  of  the  flask  is 
drawn  out  by  the  aid  of  a  blowpipe  to  a  narrow  opening.  The 
flask  is  filled  almost  to  the  narrowing  of  the  neck  with  hydro- 
chloric acid  which  is  free  from  chlorine  or  iron  chloride,  and  a 
little  sodium  carbonate  is  added  in  order  to  displace  the  last  trace 
of  air  by  carbon  dioxide.  The  flask  is  then  closed  by  sealing  off 
the  neck  in  the  blowpipe  and  warmed  in  a  water  bath  until  the 
cerium  compound  is  completely  dissolved,  and  the  quantity  of 
iodine  set  free  is  determined  by  iodometric  analysis. 

Browning  (with  Hanford  and  Hall)f  has  shown  that  good 
results  may  be  obtained  by  digesting  cerium  dioxide  with  potas- 
sium iodide  and  hydrochloric  acid  in  a  glass-stoppered  bottle 
and  determining  by  sodium  thiosulphate  the  iodine  set  free,  or 
by  distilling  and  estimating  the  iodine  passing  to  the  receiver  , 
in  accordance  with  the  reaction: 

8HCl+2KI  =  2  CeCl3  +  2  KC1  +  4  H2O  +  I2. 


Digestion  According  to  the  first  procedure,  weighed  portions 

Process.  of  the  pure  cerium  dioxide  are  placed  in  small  glass- 

stoppered  bottles  of  about  100  cm.3  capacity,  together  with 
I  grm.  of  potassium  iodide  free  from  iodate,  and  a  few  drops  of 
water  to  dissolve  the  iodide.  A  current  of  carbon  dioxide  is 
passed  into  the  bottle  for  about  five  minutes  to  expel  the  air, 
10  cm.3  of  pure  strong  hydrochloric  acid  are  added,  the  stopper 
is  inserted  and  the  bottle  heated  gently,  upon  a  steam  radiator, 
for  about  one  hour,  until  the  dioxide  dissolves  completely  and 
the  iodine  is  set  free.  After  cooling  the  bottle,  to  prevent  loss 
'of  iodine  upon  removing  the  stopper,  the  contents  are  carefully 
washed  into  about  400  cm.3  of  water  and  titrated  with  standard 

*  Ann.  Chem.,  cv,  49. 

t  Philip  E.  Browning,  with  G.  A.  Hanford  and  F.  J.  Hall,  Am.  Jour.  Sci., 
[4],  viii,  452. 


CERIUM 


247 


sodium  thiosulphate  to  determine  the  amount  of  iodine  liberated. 
Results  of  this  procedure,  corrected  by  0.04  cm.3  of  n/io  iodine, 
which  is  the  amount  of  iodine  set  free  in  blank  determinations, 
are  given  in  the  accompanying  table. 

Digestion  Process. 


CeO2  taken, 
grm. 

CeO2  found. 

grin. 

Error, 
grm. 

O.IOOO 

0.0994 

—  0.0006 

O.  1032 

0.1034 

+O.OOO2 

o.  1016 

o.  1017 

+O.OOOI 

0.1054 

o.  1041 

—  0.0013 

O.2OIO 

O.2O2I 

+O.OOII 

o.  1104 

O.IIO9 

-j-o.0005 

o.  1914 

0.1907 

—0.0007 

o.  1604 

0.1603 

—  O.OOOI 

0.2146 

0.2145 

—  O.OOOI 

o.  1108 

0.1099 

—0.0009 

0.1346 

0-1347 

+O.OOOI 

0.1540 

0-1534 

—  0.0006 

0.1976 

0.1968 

—0.0008 

0.1230 

o  .  i  240 

+0.0010 

0.1199 

O.  I2OI 

+0.0003 

0.1524 

0.1528 

+0.0004 

O.  1212 

O.I2II 

—  O.OOOI 

0.1528 

0.1543 

+0.0015 

Distillation  According  to  a  second  procedure,  portions  of  ce- 

Process.  rium  dioxide  were  weighed  out  into  the  retort  of  a 

distillation  apparatus*  consisting  of  a  Voit  flask,  serving  as  the 
retort,  sealed  to  the  inlet  tube  of  a  Drexel  wash  bottle,  used  as 
a  receiver,  the  outlet  tube  of  which  was  trapped  by  sealing  on 
Will  and  Varrentrapp  absorption  bulbs.  In  the  retort  are  placed 
the  cerium  dioxide,  15  cm.3  of  water,  I  grm.  of  potassium  iodide, 
and  10  cm.3  of  pure  strong  hydrochloric  acid.  In  the  receiver 
are  100  cm.3  of  water  and  2  grm.  to  3  grm.  of  potassium  iodide, 
and  in  the  bulbs  a  dilute  solution  of  potassium  iodide.  Before 
adding  the  hydrochloric  acid  a  current  of  carbon  dioxide  is 
passed  through  the  apparatus  for  some  minutes.  After  adding 
the  acid,  the  liquid  is  boiled  in  the  current  of  carbon  dioxide  to 
a  volume  of  15  cm.3,  and  when  the  free  iodine  has  almost  com- 
pletely left  the  retort  and  passed  into  the  receiver  the  apparatus 
is  allowed  to  cool.  The  iodine  in  the  receiver  is  titrated  directly 
*  See  Fig.  3,  page  4. 


248 


METHODS  IN  CHEMICAL  ANALYSIS 


with  sodium  thiosulphate,  and  that  in  the  retort  after  dilution 
of  the  residue  to  about  400  crn.3,  the  amount  in  the  retort  seldom 
exceeding  the  equivalent  of  a  few  drops  of  n/io  iodine  solution. 
Results  are  given  in  the  table. 


Distillation  Process. 


CeO2  taken, 
grm. 

CeO2  found, 
grm. 

Error, 
grm. 

o.  1028 

O.2O6O 

0.1013 
0.2055 

—  0.0015 
—  0.0005 

o.  2014 
o.  1716 

0.2012 

o.  1711 

—  0.0002 

—0.0005 

0.0974 

0.1600 
0.1268 
0.1276 
o.  1620 

0.0972 
0.1587 
0.1254 
0.1268 

o.  1612 

—  O.O002 
-0.0013 
—  O.OOI4 
—  O.OOOS 
—  O.OOOS 

o.  1016 

0.1548 

O.  IOII 
0.1543 

—  O.OOO5 
—  0.0005 

0.1352 

0.1342 

—  o.ooio 

The  Estimation  of  Cerium  Oxalate  by  Potassium 
Permanganate. 

Stolba  *  has  stated  that  cerium  oxalate  may  be  estimated  volu- 
metrically  in  the  same  manner  as  calcium  oxalate  by  treating  the 
washed  precipitate,  suspended  in  warm  water,  to  which  a  mod- 
erate amount  of  sulphuric  acid  has  been  added,  with  potassium 
permanganate;  that  as  the  titration  proceeds  the  precipitate 
disappears  and  the  end  reaction  is  sharp;  and  that  the  perman- 
ganate does  not  oxidize  the  cerium  from  the  lower  to  the  higher 
condition. 

Browning  and  Lynch  f  have  presented  experimental  proof  of 
the  correctness  of  this  statement,  and  have  shown  that  the  cerium 
may  be  determined  either  by  titration  of  the  precipitated  cerium 
oxalate  or  by  titration  of  the  excess  of  ammonium  oxalate  left 
over  in  the  precipitation. 

To  the  cerium  salt  in  solution  in  100  cm.3  to  200  cm.3  of  water 
is  added  a  definite  amount  of  a  standardized  solution  of  ammo- 
nium oxalate  in  excess;  the  whole  is  warmed  to  induce  a  crys- 

*  Sitzungsber.  d.  kgl.  bohm.  Gesellsch.  d.  Wissenschaften  v.  4  Juli,  1879; 
Zeit.  anal.  Chem.,  xix,  194. 

t  Philip  E.  Browning  and  Leo  A.  Lynch,  Am.  Jour.  Sci.,  [4],  viii,  457. 


CERIUM 


249 


talline  condition  of.  the  precipitate.  The  precipitate  is  filtered 
off  on  paper,  carefully  washed,  and  dissolved  by  passing  10  cm.3 
of  hot  [1:3]  sulphuric  acid  repeatedly  through  the  filter.  The 
filtrate  and  washings  are  made  up  to  about  500  cm.3  and  warmed 
to  70°  or  80°,  and  the  oxalic  acid  in  solution  is  titrated  with  per- 
manganate. The  filtrate  from  the  cerium  oxalate,  containing 
the  excess  of  ammonium  oxalate,  is  diluted  to  500  cm.3,  acidified 
with  10  cm.3  of  dilute  [i  :  3]  sulphuric  acid,  I  grm.  of  manganous 
sulphate  is  added  to  insure  regularity  of  action  in  presence  of 
hydrochloric  acid,  and  the  oxalic  acid  is  titrated  by  perman- 
ganate. 

Results  obtained  in  this  manner  are  given  in  the  table. 

Permanganate  Titration  of  Precipitate  and  of  Excess  of  Precipitant. 


Amount  taken, 
calculated  as  CeCl3. 

grm. 

Treatment  of  precipitate. 

Treatment  of  filtrate. 

Amount  found, 
calculated  as  CeCl3. 

grm. 

Error,  calculated 
as  CeCl3. 

grm. 

Amount  found, 
calculated  as 
CeCl3. 
grm. 

Error,  calcu- 
lated as 
CeCl3. 
grm. 

Precipitation  in  neutral  solution. 


o.  1091 

o.  1103 

-{-O.OOI2 

0.1091 

0.1087 

—  O.OOO4 

O.IO87 

—  0.0004 

0.1364 

0.1373 

+o  .  0009 

O.I39I 

+O.OO27 

0.1364 

0.1367 

+0.0003 

0.1367 

+0.0003 

0.2182 

O.22O2 

+O  .  OO2O 

0.2206 

+O.OO24 

Precipitation  in  acid  solution. 


0.1091 

o  1087 

—  o  0004 

0.1519 
0.1364 

0.2182 

0.1535 
0.1367 
0.2183 

+0.0016 
+0.0003 

+O.OOOI 

0.1535 
0.1367 
0.2183 

+0.0016 
+o  .  0003 

+O.OOOO 

The  Estimation  of  Cerium  in  the  Presence  of  Other  Rare  Earths  by 

the  Action  of  Potassium  Ferricyanide  in  Alkaline  Solution 

and  Potassium  Permanganate  in  Acid  Solution. 

Browning  and  Palmer*  have  shown  that  the  oxidation  of  ce- 
rium from  the  cerous  to  the  eerie  condition  may  be  effected  by 
potassium  ferricyanide  in  alkaline  solution,  registered  in  the 

*  Philip  E.  Browning  and  Howard  E.  Palmer,  Am.  Jour.  Sci.,  [4],  xxvi,  83. 


25° 


METHODS  IN  CHEMICAL  ANALYSIS 


amount  of  potassium  ferrocyanide  formed  according  to  the  follow- 
ing equation, 

2  K3FeC6N6  +  Ce203  +  2  KOH  =  2  K4FeC6N6  +  H2O  -f  2  CeO2, 

and  subsequently  determined  by  titration  with  potassium  per- 
manganate in  acid  solution. 

The  procedure  is  as  follows :  To  the  cerous  sulphate  in  solution 
20  cm.3  of  a  solution  containing  20  grm.  of  potassium  ferricyanide 
to  the  liter  are  added,  and  potassium  hydroxide  to  complete 
precipitation.  The  precipitated  hydroxide  is  filtered  off,  and 
the  filtrate  and  washings,  amounting  in  volume  to  from  200  cm.3 
to  250  cm.3,  after  being  made  distinctly  acid  with  dilute  sul- 
phuric acid,  are  titrated  with  a  standard  solution  of  potassium 
permanganate  until  the  presence  of  the  permanganate  color  shows 
the  oxidation  of  all  the  ferrocyanide  to  ferricyanide*  according 
to  the  equation : 

5  K4FeC6N6  +  KMnO4  +  4  H2SO4  =  5  KsFeCeNe  +  3  K2S04  -f- 
MnSO4  +  4  H2O. 

From  this  equation  and  the  preceding  one  the  amount  of 
cerium  present  may  be  readily  calculated. 

Each  day  before  the  ferricyanide  is  used  a  portion  of  20  cm.3 
of  the  solution  is  acidified  and  titrated  with  the  permanganate, 
and  the  correction  indicated,  generally  from  one  to  three  drops, 
is  subtracted  from  the  amount  of  the  permanganate  used  in 
actual  determinations. 

Oxidation  by  Ferricyanide  and  Titration  of  the  Reduced  Ferricyanide. 


Ce  taken,  calcu- 
lated as  Ce2O3. 

grm. 

Ce  found,  calcu- 
lated as  Ce2Os. 

grm. 

Error, 
grm. 

0.1834 

0.1819 

—  0.0015 

0.1376 

0.1380 

+0.0004 

0.1834 

o.  1829 

—  0.0005 

0.1834 

o.  1829 

—  0.0005 

0.1834 

0.1834 

o.oooo 

0.1376 

0.1385 

+o  .  0009 

0.1376 

0.1371 

-0.0005 

0.1376 

0-1374 

—  O.OO02 

0.1376 

0.1380 

+0.0004 

0.1834 

0.1824 

—  o.ooio 

0.1326 

0.1335 

+0.0009 

0.1326 

0.1328 

+O.OO02 

*  Button's  Vol.  Anal.  Qth  ed.,  page  209. 


TIN 


25* 


Oxidation  by  Ferricyanide  and  Titration  of  the  Reduced  Ferricyanide. 


Ce  taken,  calcu- 
lated as  Ce2O3. 

grm. 

Ce  found,  calcu- 
lated as  Ce2O3. 

grm. 

Error, 
grm. 

Other  rare  earths  present, 
calculated  as  oxides. 

grm. 

0.1328 

0.1335 

+0.0007 

o.    ThO2. 

0.1327 

0.1322 

—0.0005 

o.    ThO2. 

O.O266 

0.0275 

+0  .  OOOQ 

o.    ThO2. 

0.0267 

'O.O272 

-j-o  .  0005 

o.    ThO2. 

0.1324 

0.1326 

-j-O  .  OOO2 

o.    Y203. 

0.1326 

0.1323 

—0.0003 

o.    Y2O3. 

O.O266 

0.0264 

—  O.OOO2 

o.    Y203. 

0.0264 

O.O27I 

+0.0005 

o.i  Y203. 

0.1376 

0.1370 

—0.0006 

0.15  La2O3+Di2Os. 

O.IIOI 

o.  1091 

—  O.OOIO 

0.15  La203+Di2O3. 

0.1324 

0.1332 

+0.0008 

0.03  ZrO2. 

All  the  various  operations  in  this  process  are  carried  on  with- 
out warming  the  solution.  The  filtrations  and  washings  are 
made  under  gentle  pressure,  and  require  on  an  average  not  more 
than  fifteen  to  thirty  minutes.  In  the  preceding  tables  are 
results  obtained  with  cerium  alone  and  in  presence  of  sails  of 
other  rare  earths. 

This  method  presents  no  difficulties  in  manipulation  and  is 
especially  adapted  to  the  rapid  estimation  of  cerium  in  rare  earth 
mixtures. 

TIN. 

The  Electrolytic  Determination  of  Tin. 

From  a  solution  of  stannous  ammonium  chloride  in  a  satu- 
rated solution  of  ammonium  oxalate,*  Medway  has  precipitated 
the  tin  successfully  upon  the  rotating  crucible. f 

Deposition  of  Tin  on  the  Rotating  Cathode. 


Tin  taken. 

Tin  found. 

Error. 

Current. 

Time. 

N.  D.100. 

grm. 

grm. 

grm. 

Amp. 

mm. 

O  .  0804 

o  .  0802 

—  O.OOO2 

2-5 

8-3 

20 

o  .  0804 

o  .  0800 

—  O.OOO4 

2 

6.6 

20 

o.  1607 

o.  1610 

+0.0003 

2-5 

8-3 

2O 

0.  1607 

0.1603 

—  0.0004 

2-5 

8-3 

2O 

0.1607 

o.  1607 

o.oooo 

3-5 

u.  6 

15 

*  H.  E.  Medway,  Am.  Jour.  Sci.,  [4],  xviii,  56. 
f  See  Fig.  13,  page  12. 


252  METHODS  IN  CHEMICAL  ANALYSIS 

LEAD. 
The  Detection  of  Lead. 

It  has  been  shown  by  Browning  and  Blumenthal*  that  lead 
may  be  separated  as  the  sulphate,  in  association  with  the  alkali 
earth  sulphates,  and  tested  for  in  the  solution  obtained,  by  treat- 
ment of  these  insoluble  sulphates  with  ammonium  acetate,  f 

The  Electrolytic  Determination  of  Lead  as  the  Dioxide. 

In  depositing  lead  dioxide  electrolytically,  solutions  contain- 
ing nitric  acid  are  employed;  precautions  must  be  taken  in 
regard  to  concentration  of  acid,  strength  of  current  and  tem- 
perature; and  the  liquid  is  siphoned  off  before  interruption  of 
the  current. t  With  the  rotating  cathode  making  600  revolutions 
a  minute  and  a  sand-blasted  platinum  dish  for  the  anode,  Exner§ 
obtained  in  ten  to  fifteen  minutes  adherent  deposits  with  a 
current  N.  D.ioo  =  10  amp.  and  4.5  volts  acting  upon  125  cm.3  of 
solution  containing  20  cm.3  of  concentrated  nitric  acid. 

To  obviate  the  necessity  of  large  and  expensive  apparatus  of 
platinum,  Gooch  and  Beyer  ||  have  experimented  with  the  filter- 
ing crucible  used  as  a  cathode  in  the  devices  previously  described.  *[f 

In  preliminary  trials  of  electrolysis  in  the  closed  cell  with  sub- 
sequent filtration**  it  was  found  that  when  the  concentration  of 
nitric  acid  amounted  to  10  cm.3  in  60  cm.3  of  liquid,  with  a  current 
of  4  amperes  (N.  D.ioo  =  10  amp.)  and  6  volts,  two  sources  of 
error  appeared.  In  the  first  place,  the  deposition  of  metallic  lead 
upon  the  cathode  was  often  noticeable;  and  secondly,  it  appeared 
to  be  impossible  to  make  the  precipitation  of  lead  dioxide  com- 
plete so  long  as  that  substance  was  allowed  to  float  in  the  liquid. 
Similar  results  were  obtained  in  experiments  in  which  urea  was 
added  to  the  liquid  for  the  purpose  of  obviating  the  solvent  action 

*  Philip  E.  Browning  and  Philip  L.  Blumenthal,  Am.  Jour.  Sci.,  [4],  xxxii,' 
246. 

t  See  page  442. 

J  Smith,  Electro-analysis,  page  105,  edition  of  1911. 

§  Jour.  Am.  Chem.  Soc.,  xxv,  904. 

||  F.  A.  Gooch  and  F.  B.  Beyer,  Am.  Jour.  Sci.,  [4],  xxvii,  59. 

^  See  pages  13  to  20. 

**  See  Fig.  15,  page  15. 


LEAD  253 

of  dissolved  oxides  of  nitrogen  upon  lead  dioxide.  In  the  experi- 
ments with  this  form  of  apparatus,  the  stirring  of  the  asbestos 
felt  by  gas  evolved  upon  the  bottom  of  the  crucible  used  as  an 
anode,  as  well  as  the  deposition  of  oxide  on  the  outer  surface  of 
the  crucible,  was  prevented  by  taking  the  precaution  to  moisten 
the  asbestos,  from  the  outside,  with  a  drop  of  nitrobenzene  which, 
being  insoluble  in  water,  prevents  the  contact  of  the  aqueous 
electrolyte  with  the  electrode  surface  underneath  the  asbestos. 
An  increase  of  nitric  acid  to  the  proportion  of  30  cm.3  in  100  cm.3 
of  solution  served  to  prevent  the  deposition  of  lead  upon  the 
cathode,  but  to  prevent  the  re-solution  of  lead  dioxide  it  was 
found  to  be  necessary  to  use  the  process  of  continuous  filtration, 
so  that  the  deposit  might  be  compacted  upon  the  felt,  and  after 
deposition  was  complete  to  replace  the  acid  liquid  by  a  solution 
of  ammonium  nitrate  without  interruption  of  the  current.  After 
washing  out  the  nitric  acid  with  the  solution  of  ammonium 
nitrate,  the  final  washing  was  completed  with  water.  The  form 
of  apparatus  employed,  shown  in  Figure  1 6,  and  the  manner  of 
using  are  described  on  page  17.  In  the  table  are  given  the 
results  of  experiments  following  this  procedure,  and,  for  com- 
parison, the  result  of  an  experiment  in  which  it  was  found  that, 
though  electrolysis  was  continued  by  the  circulating  process  until 
the  filtrate  contained  no  lead,  traces  of  lead  dioxide  went  into 
solution  after  the  current  had  been  diminished  by  the  gradual 
dilution  with  water  used  to  replace  the  electrolyte.  Tests  for 
lead  in  filtrates  and  washings  were  made  by  neutralizing  with 
ammonium  hydroxide  and  adding  ammonium  sulphide,  or  acetic 
acid  and  potassium  chromate. 

From  the  results  of  the  experiments  described,  it  appears  that 
good  analytical  results  in  the  deposition  of  lead  dioxide  may  be 
obtained  with  the  filtering  crucible  used  as  an  electrolytic  cell  if 
nitric  acid  be  present  to  the  proportion  of  30  cm.3  of  the  con- 
centrated acid  in  100  cm.3  of  solution,  the  liquid  kept  in  con- 
tinuous filtration  until  the  electrolysis  of  the  lead  salt  is  complete, 
the  acidic  liquid  replaced  by  a  solution  of  ammonium  nitrate 
so  that  the  electric  current  passing  shall  not  fall  off  until  the  nitric 
acid  has  been  removed,  the  final  washings  made  with  water,  and 
the  deposit  weighed  after  drying  at  200°.  The  time  required 
for  the  complete  deposition  of  0.15  grm.  of  lead  dioxide  under 
the  conditions  described  is  about  two  hours. 


254 


METHODS  IN  CHEMICAL  ANALYSIS 


Electrolysis  with  Continuous  Filtration. 


Pb(NOs),. 

grm. 

Volume. 
cm.3 

HN03 
cone. 

cm.8 

Current. 

Time, 
min. 

PbO2 

found. 

grm. 

Theory 
for  PbO2. 

grm. 

Error, 
grm. 

Amp. 

N.D.ux,. 

Volt. 

A.    With  no  ammonium  nitrate  in  electrolyte  or  in  wash-water. 


0.2023 

50 

15 

2 

4 

5 

IO 

4 
5 

5 
130 

0.1460 

0.1436 

—0.0024 

B.    With  ammonium  nitrate  in  electrolyte  and  in  wash- water. 


O.2O22 

50 

15 

2 

5 

4 

40 

4 

.10 

5 

IOO 

o-i459 

o.  1462 

-f  0.0003 

0.2014 

5° 

IS 

2 

5 

4 

S 

4 

IO 

5 

US 

Q-I4S4 

o.  1458 

+o  .  0004 

O.2OOI 

SO 

15 

2 

S 

4 

5 

4 

10 

5 

n5 

0.1444 

0.1442 

—  O.OO02 

o  .  2006 

50 

15 

2 

5 

4 

5 

4 

10 

S 

H5 

o.  1448 

0.1446 

—  O.OOO2 

o  .  2046 

50 

15 

2 

S 

4 

5 

4 

IO 

S 

US 

0.1477 

0.1472 

—  0.0003 

C.    With  ammonium  nitrate  in  wash-water  only. 


O.2O2O 

50 

15 

2 

5 

4 

5 

4 

IO 

5 

H5 

0.1458 

0.1460 

+o  .  0003 

0.2037 

50 

15 

2 

5 

4 

5 

4 

4 

IO 

5 

H5 

0.1470 

0-1473 

+o  .  0003 

The  Estimation  of  Lead  by  Precipitation  as  Oxalate  and  Titration 
with  Potassium  Permanganate. 

Many  investigators  have  made  use  of  precipitation  as  the 
oxalate  for  the  estimation  of  lead.  Ward*  has  shown  that  the 
addition  of  considerable  amounts  of  acetic  acid  favors  complete- 
ness of  precipitation,  whether  ammonium  oxalate  or  oxalic  acid 
is  used  as  the  precipitant.  Precipitation  is  effected  in  the  boiling 
solution,  the  precipitated  oxalate  is  collected  on  asbestos  in  the 
perforated  crucible  and  washed  with  small  amounts  of  water. 
The  oxalic  acid  is  set  free  by  treatment  of  the  washed  precipitate 
with  warm  dilute  sulphuric  acid  and  titrated  with  permanganate. 
When  the  acetic  acid  present  does  not  exceed  one-fourth  of  the 
volume,  precipitation  is  not  quite  complete,  but  if  half  the 
volume  at  precipitation  is  made  up  of  glacial  acetic  acid,  even 
*  H.  L.  Ward,  Am.  Jour.  Sci.,  [4],  xxxiii,  334. 


LEAD 


255 


in  presence  of  moderate  amounts  of  acetates,  the  results  of  titra- 
tion  are  accurate.  The  details  of  the  preferred  treatment  are 
given  in  the  table. 

Precipitation  and  Titration  of  Lead  Oxalate. 


Lead  taken  as 
nitrate. 

grm. 

Volume  at 
precipitation. 

cm.» 

Acetic 
acid. 

cm.' 

Ammonium 
oxalate. 

grm. 

Lead  found, 
grm. 

Error, 
grm. 

Precipitation  by  ammonium  oxalate. 


0.0050 

IOO 

50 

4 

o  .  0048 

—  O.OOO2 

0.0050 

IOO 

50 

4 

0.0045 

—  0.0005 

0.0250 

IOO 

50 

4 

0.0256 

+0  .  0006 

0.0250 

IOO 

50 

4 

0.0250 

0  .  0000 

0.0500 

IOO 

50 

4 

0.0505 

+o  .  0005 

O.  IOOO 

200 

IOO 

8 

0.1002 

+O.OOO2 

Precipitation  by  oxalic  acid. 


0.0050 

50 

25 

I 

O.OO5O 

O  .  OOOO 

0.0250 

50 

25 

I 

0.0256 

+  O.OO06 

0.1000 

IOO 

50 

2 

O.  IOO2 

+0  .  OOO2 

In  presence  of  2  grm.  of  ammonium  acetate  or  potassium  acetate. 


0.1000 

IOO 

50 

2 

O.IOOO 

0  .  OOOO 

O.IOOO 

IOO 

50 

2 

0.0997 

—0.0003 

O.IOOO 

IOO 

50 

2 

O.IOOO 

o.oooo 

CHAPTER  VIII. 


NITROGEN;  PHOSPHORUS;  ARSENIC;  ANTIMONY;  BISMUTH; 

VANADIUM. 

NITROGEN. 

The  Determination  of  Nitrogen  Liberated  by  Action  of  Sodium 
Hypobromite  upon  Ammonia  Compounds  and  Derivatives. 

THE  apparatus  designed  by  Kreider*  for  the  determination  of 
volatile  products  by  loss  is  well  suited  to  the  determination  of 
the  nitrogen  liberated  from  urea,  ammonium  oxalate,  ammonium 
chloride,  etc.,  by  the  action  of  sodium  hypobromite.  Results 
obtained  by  the  use  of  this  apparatus,  described  and  figured 
elsewhere,!  are  given  in  the  accompanying  table. 
Determination  of  Nitrogen. 


Taken, 
grm. 

Found, 
grm. 

Error, 
grm. 

f 

O.IOOO 

o  .  0469 

+0.0003 

I 

O.IOOO 

0.0467 

+0.0001 

Urea               4 

O.  IOOO 

0.0467 

+O.OOOI 

O.IOOO 

o  .  0468 

+O.OOO2 

( 

O.IOOO 

0.0467 

-j-o.oooi 

O.IOOO 

0.0204 

+0.0007 

O.IOOO 

0.0197 

o.oooo 

Ammonium  oxalatc  - 

O.IOOO 

0.0198 

+O.OOOI 

O.  IOOO 

0.0198 

+O.OOOI 

O.IOOO 

0.0196 

—  O.OOOI 

O.  IOOO 

0.0264 

+O.OOO2 

O.IOOO 

0.0265 

+0.0003 

Ammonium  chloride. 

O.IOOO 

0.0261 

—  O.OOOI 

O.IOOO 

0.0263 

+O.OOOI 

O.IOOO 

0.0261 

—  O.OOOI 

The  Estimation  of  Nitrates  by  Expulsion  of  Nitrogen  Pentoxide 

on  Ignition. 

For  the  determination  of  the  nitrogen  pentoxide  combined 
in  nitrates  which  leave  definite  oxides  on  ignition,  and  estimation 
of  the  containing  nitrates,  various  fluxes  have  been  used  to  aid 
*  J.  Lehn  Kreider,  Am.  Jour.  Sci.,  [4],  xix,  188. 
t  Seepage  I. 

256 


NITROGEN 


257 


in  the  expulsion  of  the  volatile  oxide  and  to  conserve  the  residual 
oxide  in  definite  form  for  weighing.  Borax,  silicon  dioxide,  potas- 
sium dichromate  and  sodium  metaphosphate,  which  have  been 
employed  thus,  as  well  as  in  the  similar  determination  of  car- 
bon dioxide  in  carbonates*  —  all  present  certain  disadvantages. 
In  sodium  paratungstate  of  composition  corresponding  approxi- 
mately to  the  formula  5Na2O.i2WO3,  or  Nai0Wi2O4i,  Gooch  and 
Kuzirianf  find  a  material  very  easily  prepared,  stable  in  fusion, 
and  well  suited  for  use  as  a  flux  in  the  rapid  determination  of 
nitrates  by  loss  on  ignition.  The  sodium  paratungstate  is  pre- 
pared by  dehydrating  and  fusing  over  the  blast  lamp  a  known 
weight  of  normal  sodium  tungstate,  Na2WO4.2H2O,  adding  an 
equal  weight  of  tungsten  trioxide,  WO3  (previously  ignited  with 
care  to  remove  all  ammonia  and  to  insure  complete  oxidation), 
and  heating  to  clear  fusion.  The  cooled  mass,  which  is  very 
easily  pulverized,  is  ground  and  bottled.  From  this  material, 
kept  over  sulphuric  acid  (though  not  more  than  ordinarily  hygro- 
scopic), portions  are  weighed  for  the  analytical  determinations. 
Approximately  half  the  weight  of  the  paratungstate  is  tungsten 
trioxide  (molecular  weight  232),  and  this  should  be  capable  of 
expelling  nitrogen  pentoxide  (molecular  weight  108.02)  to  an 
amount  one-half  its  own  weight.  The  weights  of  paratungstate 
to  be  used  are  approximately  four  times  the  weight  of  nitrogen 
pentoxide  to  be  expelled.  It  is  a  good  practice  to  weigh  a  plati- 


Analysis  of  C.  P.  Nitrates  of  Commerce,  after  Drying. 


Nitrate  taken, 
grm. 

Na10W12041  taken, 
grm. 

Loss  on  ignition, 
grm. 

Theory  lor  N'jOj. 
grin. 

Error, 
grm. 

KN03. 

0.5000 

•5 

0   2668 

o  2670 

—  O.OOO2 

0.5000 

•5 

o  .  2678 

o  2670 

+  O.OOO8 

0.5000 

•  5 

0.2674 

o  2670 

-|-O.OOO4 

0.5000 

•  5 

0.2672 

o  2670 

+  0.0002 

0.5000 

•  S 

0.2675 

o  2670 

+  0.0005 

Sr(NO,)j. 

0.5000 

2 

0.2544 

0.2543 

+  O  OOOI 

0.5000 

3 

o  2546 

o  2543 

+0.0003 

Ba(N03)2. 

o.  5000 

3 

o  2073 

o  2067 

+0.0006 

o  5000 

3 

o  2076 

o  2067 

+o  0009 

*  See  page  226. 

f  F.  A.  Gooch  and  S.  B.  Kuzirian,  Am.  Jour.  Sci.,  [4],  xxxi,  497. 


258  METHODS  IN  CHEMICAL  ANALYSIS 

num  crucible,  introduce  the  dried  nitrate  and  weigh  again,  add 
a  suitable  amount  of  the  prepared  sodium  paratungstate,  stir 
carefully  with  a  platinum  wire  with  care  to  avoid  mechanical 
loss,  and  weigh  again.  The  crucible  is  then  heated  over  a  Bun- 
sen  burner,  first  at  very  low  heat  and  then  to  fusion  of  the  mix- 
ture for  five  minutes,  cooled  in  a  desiccator  over  sulphuric  acid, 
weighed,  and  reignited  to  test  the  constancy  of  weight.  The 
constant  weight  is  usually  got  in  the  first  ignition.  In  the  table 
are  given  the  results  of  the  estimation  of  the  nitrogen  pentoxide 
in  nitrates  according  to  the  procedure  described. 

The  Estimation  of  Nitrates  by  Reduction  with  a  Ferrous  Salt  and 
Titration  of  the  Residual  Unoxidized  Salt. 

In  the  methods  for  the  quantitative  estimation  of  nitrates 
which  depend  upon  the  reduction  in  presence  of  acid  by  a  ferrous 
salt  and  the  determination  of  the  amount  of  oxidation  produced, 
scrupulous  care  is  necessary  that  the  atmosphere  in  contact  with 
the  ferrous  salt  while  the  nitrogen  dioxide  is  present  shall  be  free 
from  oxygen.  This  fact  was  recognized  by  Fresenius,*  who 
modified  the  original  process  of  Pelouzef  by  filling  the  flask  with 
carbon  dioxide  or  hydrogen  at  the  outset.  Ederf  used  carbon 
dioxide  similarly.  Holland's  method, §  roughly  described,  con- 
sists in  boiling,  until  the  air  is  expelled,  the  solution  of  the 
nitrate  in  a  flask  provided  with  a  doubly  bent  exit  tube, 
rubber- join  ted  and  fitted  with  a  pinchcock;  then  admitting 
through  the  tube  as  the  flask  cools  a  mixture  of  ferrous  salt 
and  strong  hydrochloric  acid,  heating  the  mixture  on  a  water 
bath,  and,  finally,  titrating  the  resulting  ferric  salt  with  stan- 
nous  chloride. 

All  of  these  methods  give  high  results,  either  on  account  of  the 
oxygen  invariably  present  in  carbon  dioxide  as  produced  in  the 
laboratory,  or  because  of  slow  leakage  through  the  rubber  con- 
nections during  the  long  heating,  or  because  the  nitrogen  dioxide 
is  not  driven  out  completely  from  the  solution  of  the  iron  salts, 
and  acts  with  atmospheric  oxygen  to  oxidize  the  ferrous  salt 
during  the  titration. 

*  Zeit.  anal.  Chem.,  i,  32. 
t  Ann.  Chim.,  [3],  xx,  120. 
t  Zeit.  anal.  Chem.,  xvi,  267. 
§  Chem.  News,  xvii,  219. 


NITROGEN  259 

Phelps*  has  shown,  however,  that  the  oxidation  of  ferrous 
sulphate  by  a  nitrate  in  presence  of  fairly  strong  hydrochloric 
acid  may  be  accomplished  quantitatively  by  the  aid  of  the  ap- 
paratus to  be  described,  the  standard  of  the  solution  of  the 
ferrous  salt  and  the  amount  remaining  after  the  action  being 
determined  either  iodometrically  or  by  titration  with  potassium 
permanganate . 

This  apparatus  consists  of  a  25O-cm.3  boiling  flask,  closed  with 
a  rubber  stopper  carrying  a  stoppered  funnel  of  50  cm.3  capacity 
to  serve  as  an  inlet  tube  and  a  glass  tube  of  0.8  cm.  bore  to  serve 
as  an  outlet  tube.  The  inlet  tube  is  constricted  at  the  lower 
end,  and  the  outlet  tube  is  enlarged  just  above  the  stopper  to  a 
small  bulb  (to  prevent  mechanical  loss  during  the  boiling)  and 
bent  twice  at  right  angles.  The  flask  is  supported  above  a 
Bunsen  burner  and  the  outlet  tube  dips  under  mercury  contained 
in  a  test  tube. 

The  nitrate  to  be  analyzed  is  introduced  with  water  into  the 
flask,  the  stem  of  the  separating  funnel  being  left  full  of  water, 
the  outlet  tube  is  adjusted  to  just  touch  the  surface  of  the  mer- 
cury in  the  trap,  and  air  is  expelled  from  the  flask  by  boiling  the 
solution  to  small  volume.  An  amount  of  standardized  ferrous 
sulphate  solution  known  to  be  in  excess  is  introduced  into  the 
separating  funnel,  the  outlet  tube  plunged  a  centimeter  or  two 
deep  into  the  mercury  (which  is  readily  accomplished  by  chang- 
ing the  position  of  the  flask  on  the  wire  gauze,  provided  that  the 
gauze  is  depressed  well  at  the  center  and  the  flask  is  set  well  up 
on  the  higher  part  at  the  beginning  of  the  operation),  and  then 
the  flame  withdrawn  until  diminution  of  pressure  sufficient  to 
draw  the  ferrous  solution  into  the  flask  is  made  evident  by  the 
rise  of  the  mercury  in  the  outlet  tube.  By  applying  and  with- 
drawing the  flame  and  by  regulating  the  rate  of  inflow  of  the 
solution,  the  ferrous  salt  may  be  introduced  without  admitting 
air,  and  the  funnel  washed  carefully  with  an  amount  of  con- 
centrated hydrochloric  acid  nearly  enough  to  equal  the  total 
volume  of  the  liquid  in  the  flask.  After  the  pressure  has  been 
restored  in  the  apparatus  by  heating  the  flask,  the  exit  tube  is 
again  raised  to  the  surface  of  the  mercury  and  the  solution  in 
the  flask  boiled  to  a  volume  of  10  cm.3  to  15  cm.3.  The  excess  of 
acid  is  then  nearly  neutralized  by  introducing  sodium  carbonate 
*  I.  K.  Phelps,  Am.  Jour.  Sci.,  [4],  xiv,  440. 


260 


METHODS  IN  CHEMICAL  ANALYSIS 


in  solution,  the  carbon  dioxide  evolved  assisting  in  maintaining 
the  pressure  in  the  apparatus  so  that  the  condensed  liquid  in  the 
test  tube  which  may  contain  oxidized  nitrogen  dioxide  shall  not 
be  returned  to  the  iron  solution.  The  flask  is  cooled  and  the 
ferrous  salt  remaining  determined  by  titration  with  potassium 
permanganate  after  dilution  with  600  cm.3  of  water  and  addition 
of  2  grm.  to  3  grm.  of  crystallized  manganous  chloride,  or  iodo- 
metrically  by  first  introducing  Rochelle  salt  (3  grm.)  in  solution 
and  then  neutralizing  with  acid  potassium  carbonate,  after  which 
are  added  in  succession  a  saturated  solution  of  acid  potassium 
carbonate,  iodine,  starch  paste,  standard  arsenious  acid  solu- 
tion to  the  bleaching  of  the  starch  blue,  and,  finally,  iodine  to 
coloration. 

It  was  shown  experimentally  that  prolonged  boiling  after  the 
dark  compound  of  nitrogen  dioxide  with  the  ferrous  salt  is 
broken  up  is  essential  and  that  ammonium  salts  must  be  absent 
if  the  highest  accuracy  is  desired.  Experimental  tests  of  the 
method  as  outlined  are  given  in  the  table. 

Reduction  of  Nitrate  by  Ferrous  Sulphate:    Titration  of  Excess. 


KNO3 

taken. 

Oxygen  value 
of  ferrous  salt 
taken. 

Oxygen  value 
of  ferrous  salt 
found. 

Error  on 
oxygen. 

KN03  found. 

Error  on 
KNO3. 

grm. 

grm. 

grm. 

grm. 

grm.       , 

grm. 

0.0500 

0.01823 

0.00621 

+0.00015 

o  .  0506 

+0  .  0006 

0.0500 

0.01865 

0.00681 

—0.00003 

0.0499 

—  o.oooi 

0.0500 

0.01954 

0.00768 

o.ooooo 

O.O50O 

o.oooo 

O.  IOOO 

0.02881 

0.00507 

+0.00001 

O.  IOOO 

o.oooo 

O.  IOOO 

0.02822 

0.00441 

+0.00008 

0.1003 

+0.0003 

0.2500 

0.06453 

0.00512 

+0.00009 

0.2503 

+0.0003 

0.5000 

0.13394 

0.01524 

+0.00005 

0.5002 

+  O.OOO2 

0.5000 

O.I22IO 

0.00340 

+0.00005 

0.5002 

+O.OOO2 

The  Estimation  of  Nitrates  by  Reduction  with  Ferrous  Chloride 
and  Measurement  of  the  Nitrogen  Dioxide  Evolved. 

In  consequence  of  the  fact  that  analytical  methods  in  which 
nitrates  are  estimated  by  the  amount  of  nitrogen  dioxide  evolved 
in  the  reaction  with  ferrous  salts  in  presence  of  acid  give  un- 
expectedly low  results,  the  form  of  apparatus  to  be  used  in 
such  processes  and  the  conditions  affecting  accuracy  have  been 
made  the  object  of  study  by  Roberts.*  As  the  result  of  such 
*  Charlotte  F.  Roberts,  Am.  Jour.  Sci.,  [3],  xlvi,  126. 


NITROGEN  261 

study,  it  is  pointed  out  that  to  insure  rapid  action,  the  hydro- 
chloric acid  used  should  be  fairly  strong,  and  that  in  the  measure- 
ment of  the  volume  of  nitrogen  dioxide  (NO),  swept  along  by 
carbon  dioxide  and  collected  over  sodium  hydroxide,  the  best 
analytical  results  are  obtained  when  the  gas  is  passed  through  a 
solution  of  potassium  iodide  to  break  up,  before  the  measurement 
is  made,  any  higher,  soluble  oxides  of  nitrogen  which  may  have 
been  formed  in  the  process,  notably  when  the  ferrous  salt  is 
present  in  only  small  excess  and  when  the  reaction  of  decomposi- 
tion takes  place  in  the  hot  solution.  Note  is  made  of  the  fact 
that  the  action  of  traces  of  intermixed  air  upon  the  nitrogen 
dioxide  is  to  form  nitrogen  trioxide  which  is  soluble,  and  that  the 
nitrogen  of  the  air  is  added  in  exactly  the  proportion  according 
to  which  the  nitrogen  dioxide  is  removed  with  the  disappearing 
oxygen.  When,  however,  the  gas  carrying  nitrogen  trioxide 
meets  with  potassium  iodide  before  measurement,  the  original 
amount  of  nitrogen  dioxide  is  regenerated.  Nevertheless,  ex- 
perience shows  that  the  danger  of  error  introduced  by  the  use 
'of  potassium  iodide  when  traces  of  air  are  present  is  small  in 
comparison  with  the  danger  of  error  due  to  the  presence  of 
higher  oxides  of  nitrogen  produced  in  the  main  reaction.  More- 
over, there  is  always  a  small  counterbalancing  error  due  to  the 
solubility  of  nitrogen  dioxide  in  the  solution  of  sodium  hydroxide 
over  which  it  is  measured. 

The  apparatus  found  to  be  most  satisfactory  for  this  work 
consists  of  a  small  tubulated  retort,  upon  the  neck  of  which  is 
fitted  a  small  condenser  to  prevent  loss  of  liquid  during  the  dis- 
tillation. Into  the  tubulature  of  this  retort  is  fitted  tightly,  by 
a  carefully  ground  joint,  a  tube  drawn  out  so  as  to  dip  below 
the  surface  of  the  liquid,  and  fitted  with  carefully  ground  stop- 
cocks, —  as  shown  in  the  figure,  —  and  so  branched  above  as 
to  make  it  possible  to  transmit  carbon  dioxide  through  the  ap- 
paratus, or  to  admit  any  liquid  without  introducing  air.  The 
condenser  is  joined  to  a  Will  and  Varrentrapp  bulb  contain- 
ing a  solution  of  potassium  iodide  as  a  trap,  and  this  in  turn  is 
connected  by  thick  vacuum  tubing  with  a  Hempel  gas  burette 
charged  with  a  strong  solution  of  sodium  hydroxide.  Carbon 
dioxide  is  generated  in  a  Kipp's  apparatus  by  action  of  boiled 
hydrochloric  acid  upon  boiled  marble,  and  the  liquid  is  charged 
with  cuprous  chloride,  following  Warrington's  device,  to  take  up 


262 


METHODS  IN   CHEMICAL  ANALYSIS 


traces  of  dissolved  oxygen.  Notwithstanding  all  precautions, 
however,  the  gas  from  the  generator  is  never  so  pure  that  a 
hundred  cubic  centimeters  of  it  will  not  leave  a  tiny  bubble  when 
shaken  with  a  solution  of  caustic  soda. 

In  using  this  apparatus,  the  nitrate  (about  o.i  grin,  of  potas- 
sium nitrate)  is  introduced  into  the  retort,  generally  in  the  dry 
condition,  carbon  dioxide  is  passed  through  the  apparatus  until 
the  gas  collected  over  sodium  hydroxide  leaves  only  the  minute 
bubble  which  the  gas  from  the  generator  alone  has  been  found  to 


Fig.  21. 

give,  and  40  cubic  centimeters  of  a  boiled  solution  of  ferrous 
chloride  in  hydrochloric  acid  are  admitted  through  the  funnel  tube, 
after  shutting  off  the  carbon  dioxide  and  lowering  the  leveling 
tube  of  the  Hempel  burette.  With  the  stopcocks  arranged  as 
in  sketch,  the  liquid  is  then  slowly  heated  to  boiling  and  the 
process,  continued  until  the  reaction  of  the  ferrous  salt  upon 
the  nitrate  is  apparently  complete,  when  the  carbon  dioxide  is 
again  passed  through  the  apparatus  to  secure  complete  removal 
of  the  nitrogen  dioxide,  the  absorption  of  the  carbon  dioxide  being 
hastened  by  inclining  and  shaking  the  burette  at  intervals.  The 
volume  of  the  gas  under  existing  barometric  and  thermometric 


NITROGEN 


263 


conditions  is  noted  and  from  this  the  weight  of  the  nitrate  may 
be  calculated.  Results  obtained  in  this  manner  are  given  in 
the  table. 

Reduction  by  Ferrous  Chloride:   Measurement  of  Nitrogen  Dioxide. 


KNO,  taken, 
grtn. 

KNO3  found, 
grin. 

Error, 
grm. 

O.IOOO 

0.0990 

—  0.0010 

O.IOOO 

0.1005 

+0.0005 

O.IOOO 

0.0992 

—0.0008 

O.IOOO 

0.0994 

—0.0006 

O.IOOO 

0.1008 

+0.0008 

O.IOOO 

o  .  0989 

—  O.OOII 

Action  of 
Manganous 
Chloride  in 
Hydrochloric 
Acid. 


The  lodometric  Determination  of  Nitrates. 

Noting  that  a  saturated  solution  of  manganous 
chloride  in  concentrated  hydrochloric  acid  acts  like 
the  solution  of  ferrous  chloride  in  hydrochloric  acid 
in  inducing  the  easy  decomposition  of  nitrates,  with 
the  difference,  however,  that  all  products  of  oxidation  may  be 
distilled,  while  the  metal  chloride  reverts  to  its  original  form, 
Gooch  and  Gruener  *  have  applied  this  reagent  to  the  quantita- 
tive estimation  of  nitric  acid  as  well  as  in  the  qualitative  test. 

The  solution  of  manganous  chloride  in  hydrochloric  acid  acts 
but  slowly  upon  nitrates  at  the  ordinary  temperature,  but  upon 
warming  the  decomposition  of  the  nitrate  begins  at  once  with  the 
formation  of  a  higher  chloride  of  manganese  and  liberation  of 
nitrogen  dioxide.  Ultimately,  if  heating  is  continued,  chlorine 
of  the  higher  chloride  is  evolved  and  manganous  chloride  remains. 
During  the  process  of  heating  the  color  of  the  solution  passes 
from  the  original  characteristic  green  through  darker  shades  to 
black,  and  returns  by  the  reverse  changes  to  the  original  tint. 
The  decomposition  of  the  nitrate  extends  under  the  conditions  to 
the  last  traces,  but  the  breaking  up  of  the  nitrates,  with  the 
formation  of  the  higher  chloride,  does  not  take  place  completely 
in  the  presence  of  water  amounting  to  more  than  a  half  of  the 
volume  of  strong  acid,  and  an  action  already  established  in 
strong  acid  is  reversed  by  the  addition  of  a  large  amount  of  water. 
Chlorates,  peroxides  and  other  substances  which  liberate  oxygen 

*  F.  A.  Gooch  and  H.  W.  Gruener,  Am.  Jour.  Sci.,  [3],  xliv,  117. 


264 


METHODS  IN  CHEMICAL  ANALYSIS 


or  chlorine  when  in  contact  with  strong  hydrochloric  acid  induce 
similar  phenomena,  but  in  the  absence  of  such  other  substances 
the  reaction  serves  to  detect  nitrates  when  present  in  fairly  small 
amounts  (perhaps  one  part  in  sixty  thousand),  as  shown  in  the 
accompanying  table : 


KNOa  taken. 

MnCl2.4H7O  in 
strong  HC1. 

Color  developed. 

grrn. 

cm.3 

O.OIOOO 

IO 

Black. 

o  .  00500 

5 

Black. 

O.OOIOO 

5 

Dark  brown. 

o  .  00050 

5 

Dark  green. 

O.OOO25 

5 

Deepened  tint. 

O.OOOI5 

5 

Deepened  tint. 

o  .  00005 

5 

None. 

O  .  OOOOO 

5 

None. 

In  applying  this  reaction  to  the  quantitative  estimation  of 
nitrates,  the  nitrate  to  be  estimated  is  treated,  in  an  atmosphere 
of  carbon  dioxide,  with  a  saturated  solution  of  crystallized 


Fig.  22. 

manganous  chloride  in  concentrated  hydrochloric  acid,  the  vola- 
tile products  of  action  —  chlorine,  nitrogen  dioxide  and  perhaps 
nitrosyl  chloride — are  passed  into  a  solution  of  potassium  iodide, 
and  the  iodine  set  free  is  titrated  by  sodium  thiosulphate.  The 


NITROGEN 


265 


operation  is  conducted  in  an  apparatus  made  wholly  of  glass 
where,  by  any  possibility,  rubber  connections  might  be  acted 
upon.  The  retort  used  was  a  pipette  bent  and  fitted  as  shown  in 
Fig.  22.  To  the  retort  are  sealed  Will  and  Varrentrapp  nitrogen 
bulbs,  the  outlet  tube  of  which  is  drawn  out  so  that  it  may  be 
pushed  well  within  the  inlet  tube  of  the  second  receiver  —  a 
Will  and  Varrentrapp  absorption  flask  —  and  held  in  place  by 
an  outside  rubber  connector.  The  third  receiver  acts  simply  as 
a  trap  to  exclude  air  from  the  absorption  apparatus  proper.  In 
conducting  the  experiment  the  receivers  were  charged  with 
solutions  of  potassium  iodide,  the  first  containing  three  grams, 
the  second  one  gram,  and  the  third  only  a  fraction  of  a  gram  for 
every  tenth  of  a  gram  of  nitrate  used.  The  first  receiver  was 
cooled  in  water  during  the  subsequent  process  of  distillation. 

Decomposition  of  Nitrate  by  Hydrochloric  Acid  and  Manganese  Chloride:   Titra- 
tion  of  Iodine  Set  Free  by  Volatile  Products. 


KNO3  taken, 
grm. 

MnCl2  mixture. 
cm.3 

KNO3  found, 
grm. 

Error  in  terms  of 
KNO3. 

grm. 

Error  in  terms  of 
HNO3. 

grm. 

O  .  2038 
0  .  2053 
o.  1032 
o.  1017 

2O 
2O 
10 
IO 

0.2047 
0.2057 
0.1035 
0.1004 

+0  .  OOOQ 

+o  .  0004 

+0.0003 
—  0.0013 

+0.0005 

+o  .  0003 

+O.OOO2 
—  O.OOOS 

o.  1049 

IO 

o.  1040 

O.OOOO 

O.OOOO 

o.  1027 
0.0524 

IO 
IO 

0.1023 
0.0526 

—0.0004 

+O.OOO2 

—0.0003 

+O.OOOI 

0.0513 

IO 

O  .  05  1  2 

—  o.oooi 

—  O.OOOI 

0.0354 

10 

0.0350 

—  0.0004 

—0.0003 

0.0232 

0.0107 
0.0127 

10 

5 
5 

O.O23O 

0.0106 
0.0130 

—  O.O002 
—  O.OOO4 
+0.0003 

—  O.OOOI 
—  O.OOOI 
+O.OOO2 

0.0145 

5 

0.0143 

—  O.0002 

—  O.OOOI 

0.0053 

o  .  0043 
0.0014 

5 
5 
5 

0.0052 
0.0047 
0.0018 

—  O.OOOI 

+o  .  0004 
-j-o  .  0004 

—  O.OOOI 

+0.0003 
+0.0003 

o.oooo 

5 

o.oooo 

O.OOOO 

o.oooo 

The  nitrate  and  the  manganous  mixture  following  it  are  intro- 
duced by  applying  gentle  suction  to  the  end  of  the  absorption 
train.  The  current  of  carbon  dioxide  is  started  immediately 
after  putting  in  the  manganous  mixture.  After  a  suitable  time 
has  elapsed  for  the  removal  of  air,  heat  is  applied  to  the  retort 
and  the  distillation  is  continued  until  nearly  all  the  liquid  has 
passed  over.  Finally,  the  contents  of  the  receivers  are  united, 


266  METHODS  IN  CHEMICAL  ANALYSIS 

the  washing  of  the  bulbs  was  effected  easily  and  expeditiously  by 
passing  the  wash-water  directly  through  retort  and  receiver,  the 
introduction  of  the  manganese  chloride  into  the  distillate  being 
not  at  all  prejudicial  to  the  accuracy  of  the  titration.  The 
estimation  of  free  iodine  is  made  by  titration  with  sodium  thio- 
sulphate  as  soon  as  may  be  after  admitting  air  to  the  distillate,, 
in  order  that  traces  of  dissolved  nitric  oxide  may  not  be  reoxi- 
dized  and  again  react  upon  the  iodide  present  to  liberate  more 
iodine.  The  results  of  the  experiments  conducted  in  this  man- 
ner are  given  in  the  table. 

Various  attempts  to  utilize  the  reduction  of  ar- 

Distillation  with 

Phosphoric  Acid  seme  acid  brought  about  upon  heating  a  mixture 
and  Potassium  of  standard  potassium  iodide,  potassium  arsenate 

Iodide  and  .  .  .  .  . 

Determination  and  sulphuric  acid  with  the  nitrate,  as  in  the  similar 
of  iodine  in  process  for  estimating  chlorates,*  have  proved  to  be 
futile.  The  decomposition  of  the  last  traces  of  ni- 
trates by  the  action  of  potassium  iodide  and  sulphuric  acid  does 
not  occur  except  at  concentrations  so  great  that  the  sulphuric 
acid  itself  liberates  iodine  from  iodides.  Gruenert  has  shown, 
however,  that  when  phosphoric  acid  is  substituted  for  sulphuric 
acid  the  iodine  evolved  in  the  distillation  of  such  a  mixture  may 
be  taken  as  the  measure  of  nitrates  present  in  small  amounts, 
provided  the  concentration  of  the  residue  is  not  carried  so  far  as 
to  bring  about  reduction  of  nitric  acid  to  ammonia!  nor  the  tem- 
perature so  raised  by  removal  of  water  and  elevation  of  the 
boiling  point  of  the  phosphoric  acid  as  to  cause  dissociation  of 
hydriodic  acid. 

In  Gruener's  experiments  a  small  retort  was  used,  the  neck  of 
which  was  bent  downward  about  two  inches  from  the  body,  so 
that  the  retort  itself  might  be  tipped  backward,  allowing  the 
unbent  portion  of  the  retort  to  run  upward,  thus  guarding  against 
loss  from  spattering.  Into  the  tubulature  of  the  retort  was 
ground  a  glass  tube  drawn  out  at  both  ends  to  serve  as  a  perfo- 
rated stopper  for  the  entrance  of  carbon  dioxide.  The  neck  was 
passed  through  a  rubber  stopper  into  a  side-neck  Erlenmeyer 
flask,  the  exit  tube  of  which  was  prolonged  and  dropped  into  a 
side-neck  test  tube  used  as  a  trap.  The  retort  was  covered  with  a 

*  See  page  463. 

t  Hippolyte  Gruener,  Am.  Jour.  Sci.,  [3],  xlvi,  42. 

\  Chapman,  Jour.  Chem.  Soc.,  xx,  166  (1867). 


NITROGEN 


267 


simply  contrived  hood  which  kept  the  upper  parts  warm  and  pre- 
vented the  iodine  from  settling  anywhere.  In  the  retort  was  placed 
the  nitrate  with  an  excess  of  potassium  iodide,  and  in  the  receiver 
a  known  amount  of  decinormal  solution  of  arsenious  oxide  strongly 
alkaline  with  hydrogen  sodium  carbonate  and  diluted  to  a  con- 
venient bulk.  The  trap  contained  nothing  but  water. 

The  method,  so  far  as  it  is  applicable,  may  be  summed  up  as 
follows:  The  nitrate,  not  to  exceed  the  equivalent  of  0.05  grm.  of 
potassium  nitrate,  is  introduced  into  the  retort,  with  ten  times 
its  weight  of  potassium  iodide,  and  17  cm.3  to  20  cm.3  of  phos- 
phoric acid,  of  specific  gravity  1.43.  All  water  used  should  be 
recently  boiled.  Carbon  dioxide  is  passed  from  a  generator  set 
up  with  materials  carefully  boiled  and  containing  cuprous 
chloride  to  take  up  the  oxygen  from  any  traces  of  air.  The  neck 
of  the  retort  passes  into  a  receiver  containing  a  known  amount 
of  decinormal  arsenious  oxide,  alkaline  with  a  good  excess  of 
hydrogen  sodium  carbonate  and  diluted  to  a  convenient  bulk. 
To  this  flask  is  attached  for  additional  safety  a  simple  trap  con- 
taining water.  The  solution  in  the  retort  is  boiled  until  it  is 
clear  that  no  more  iodine  remains,  when  the  receiver,  after  proper 
washing  and  addition  of  the  liquid  in  the  trap,  is  titrated  with 
iodine  to  find  the  amount  of  arsenious  oxide  still  left.  This  gives 
the  measure  of  the  iodine  evolved  and  consequently  of  the 
nitrate  present,  according  to  the  equation : 

2  HNO3  +  6HI  =  4H2O  +  2  NO  +  3  I2. 
The  details  of  test  determinations  are  given  in  the  table : 

Decomposition  by  Phosphoric  Acid  and  Iodide:    Estimation  of  Iodine  Set  Free* 


KNO3 

taken. 

KI  taken. 

Found. 

Specific  grav- 
ity of  solu- 
tion of 

Amount  of 
solution 
used. 

Error, 
KNO3. 

Error, 
HNO,. 

phosphoric 

acid. 

grin. 

grm. 

grm. 

cm.* 

grm. 

grm. 

0.0500 

I 

0.0500 

•43 

17 

o.oooo 

0.0000 

O.O2OO 

o-5 

O.O2OI 

•43 

17 

+0.0001 

+O.OOOI 

O.O2OO 

0.0198 

•43 

17 

—  0.0002 

—o.oooi 

0.0250 

0.0250 

•43 

17 

0.0000 

o  .  oooo 

0.0300 

0.0307 

•43 

17 

+0.0007 

+0.0004 

0.0300 

0.0312 

•43 

17 

+O.OOI2 

+0.0007 

0.0350 

0-0353 

•43 

17 

+0.0003 

+O.OOO2 

o  .  0400 

o  .  0409 

•35 

2O 

+0.0009 

+o  .  0006 

0.0450 

o  .  0444 

•35 

2O 

—0.0006 

—0.0004 

0.0500 

o  .  0499 

•37 

2O 

—  o.oooi 

—o.oooi 

268  METHODS  IN  CHEMICAL  ANALYSIS 

The  process  is  good  for  estimating  nitrates  in  quantities  not 
exceeding  the  equivalent  of  0.04  grm.  or  0.05  grm.  of  potassium 
nitrate.  With  quantities  of  nitrate  above  0.05  grm.  it  is  not  safe, 
inasmuch  as  with  a  moderate  amount  of  water  present  some  nitric 
acid  distils  over  undecomposed  and  with  little  water  present 
other  complications  arise. 

To  register  the  action  of  nitrates,  Gruener*  tried 
Decomposition    the  effect  of  antimony  trichloride  in  hydrochloric 
acid  and  showed  that  the  reaction  proceeds  mainly 


Determination    according  to  the  equation, 

of  Oxidation  in 

Residue  and          3SbCl3  +  2  HNO3  +  6HC1  =  3  SbCl5  +  2  NO  +  5  H2O, 

of  Iodine  in 

Distillate.  and  in  smaller  degree  according   to  the  equation, 

SbCl5  +  2  NO  =  SbCl3  +  2  NOC1. 

Nitrosyl  chloride  fails  to  oxidize  arsenious  oxide  in  alkaline 
solution,  breaking  up  hydrolytically  into  hydrochloric  acid  and 
nitrous  acid,  but  from  acidulated  potassium  iodide  out  of  contact 
with  air  it  sets  free  quantitatively  an  amount  of  iodine  correspond- 
ing to  the  chlorine.  These  reactions  may,  therefore,  be  ap- 
plied together  to  the  estimation  of  the  nitrate,  by  noting  both 
the  iodine  evolved  in  the  action  of  nitrosyl  chloride  upon  potas- 
sium iodide  in  the  distillate  and  the  degree  of  oxidation  of  the 
previously  standardized  antimony  salt  in  the  residue.  The  pro- 
cedure is  as  follows: 

Into  a  diminutive  retort  —  made  from  a  pipette,  shaped  like 
a  Liebig's  drier  f  and  connected  by  a  sliding  joint  covered  by 
rubber  with  a  Kjeldahl  tube  used  as  a  receiver,  and  so  placed 
that  carbon  dioxide  passing  through  the  apparatus  shall  enter 
from  below  —  the  dry  nitrate  is  introduced,  and  washed  down 
with  a  few  drops  of  recently  boiled  water,  or,  if  more  liquid 
is  required,  with  hydrochloric  acid.  From  a  burette  a  definite 
amount  of  antimonious  chloride  solution,  somewhat  in  excess  of 
the  nitrate  taken,  is  added. 

The  receiver  is  charged  with  potassium  iodide  in  recently 
boiled  water  and  is  joined  to  a  trap  filled  with  water.  After 
carbon  dioxide  has  been  passed  through  the  apparatus  for  about 
ten  minutes,  the  solution  is  warmed  in  a  bath  at  even  tempera- 
ture (iO3°-iO7°)  to  insure  the  safety  of  the  retort,  to  keep  the 

*  Am.  Jour.  Sci.,  [3],  xlvi,  47. 
t  See  Fig.  5,  page  5. 


NITROGEN 


269 


antimony  pentachloride  from  breaking  up,  to  retain  the  bulk  of 
the  acid  in  the  retort,  and  to  prevent  mechanical  loss.  After  fif- 
teen minutes'  digestion  the  contents  of  the  receiver  and  trap  are 
washed  out  and  at  once  titrated  with  sodium  thiosulphate  or 
neutralized  and  titrated  with  standard  arsenite.  The  residue  in 
the  retort  is  treated  exactly  as  was  the  antimonious  chloride 
when  it  was  standardized,  viz.,  by  dissolving  in  hydrochloric 
acid,  adding  tartaric  acid,  diluting,  nearly  neutralizing  with 
sodium  hydroxide  with  careful  cooling  to  prevent  action  of  the 
tartaric  acid  upon  antimony  pentachloride,  treating  with  an 
excess  of  acid  sodium  carbonate,  and  titrating  with  decinormal 
iodine  in  presence  of  starch.  Below  are  given  the  results  of 
experimental  tests. 

Decomposition  by  Antimony  Trichloride:   Determination  of  Oxidation  in 
Residue  and  of  Iodine  in  Distillate. 


KN03 

taken. 

grm. 

KNO3from 
SbCl5  in 
residue. 

grm. 

KNO3  from 
I  in  receiver. 

grm. 

Entire  KNO3 
found. 

grm. 

Error  in 
KNO3. 

grm. 

Error  in 
HNO3. 

grm. 

0.0222 

0.0213 

O.OO2O* 

0.0233 

+O.QOII 

+O.OOO7 

0.0336 

0.0307 

O.OO26* 

0.0333 

—  0.0003 

—  O.OOO2 

0.0470 

o  .  0436 

0.0045* 

0.0471 

+O.OOOI 

+O.OOOI 

0-0553 

0.0497 

0.0057* 

0.0554 

+0.0001 

+0.0001 

o  .  0664 

0.0673 

0.0076* 

0.0679 

+0.0015 

+0.0009 

0.0759 

0.0670 

0.0082* 

0.0752 

—0.0007 

—  0.0004 

0.0837 

0.0730 

O.OIO3* 

0.0841 

+0.0004 

+O.OOO2 

0.0934 

0.0842 

O.OII3* 

0.0955 

+O.OO2I 

+O.OOI3', 

0.1034 

O.OQO2 

0.0134* 

o.  1036 

+O.OOO2 

+O.OOOI 

O.O262 

0.0235 

O.OO24* 

0.0259 

—  0.0003 

—  O.OOO2 

O.OI27 

0.0123 

0.0007* 

0.0130 

+0.0003 

+O  .  OOO2 

0.0065 

0.0064 

0.0003* 

0.0067 

+O.OOO2 

+O.OOOI 

O.OO26 

O.OO22 

O.OOOI* 

0.0023 

—  0.0003 

—  0.0002 

0.1232 

o.  1129 

0.0098* 

0.1227 

—  0.0005 

—  0.0003 

0  .  I  540 

0.1394 

0.0146* 

0.1540 

O.OOOO 

O.OOOO- 

0.1878 

0.1655 

O.O2IO* 

0.1865 

—  0.0013 

—0.0008 

o  .  0530 

0.0481 

O.OO52t 

0.0533 

+0.0003 

+O.OOO2 

0.0547 

o  .  0484 

0.0065t 

o  .  0549 

+O.OOO2 

+  O.OOOI 

0.0541 

0.0474 

0.0063! 

0.0537 

—  0.0004 

—  O.OOO2 

Found  by  thiosulphate. 


t  Found  by  arsenite  after  neutralization. 


The  lodometric  Determination  of  Nitrites. 

The  apparatus  (consisting  of  a  boiling  flask  fitted  with  a 
stopper  which  carries  a  stoppered  funnel  and  outlet  tube  dipping 
in  a  mercury  trap)  previously  used  by  Phelps*  for  the  determina- 
tion of  nitric  acid  has  been  applied  by  him  in  the  determination 

*  See  page  259. 


270  METHODS  IN  CHEMICAL  ANALYSIS 

of  nitrites.*  In  this  method  the  nitrite  is  reduced  by  the  action 
of  potassium  iodide  and  arsenious  acid  in  acid  solution  and  meas- 
ured by  titration  of  the  arsenite  remaining,  after  neutralization. 

An  amount  of  standard  arsenite  solution,  slightly  in  excess  of 
that  required  to  take  up  the  iodine  to  be  set  free  later  by  the 
nitrous  acid,  and  25  cm.3  of  a  concentrated  solution  of  sodium 
carbonate,  are  placed  in  the  flask.  The  stem  of  the  stoppered 
funnel  is  completely  filled  with  water,  the  rubber  stopper  inserted 
tightly  and  the  contents  of  the  flask  boiled  until  all  air  is  expelled, 
a  process  requiring  an  active  boiling  of  5-8  minutes.  The  flame  is 
then  removed,  the  outlet  tube  is  plunged  deep  into  the  mercury, 
the  flask  is  cooled  with  ice  water,  and  enough  sulphuric  acid 
[1:3]  (7  cm.3)  is  sucked  in  through  the  funnel  tube  to  nearly 
decompose  the  sodium  carbonate  previously  added  and  liberate 
carbon  dioxide  to  balance  the  atmospheric  pressure.  When  the 
inward  and  outward  pressures  have  been  equalized  the  outlet 
tube  is  raised  so  that  the  end  shall  dip  in  the  water  layer  con- 
densed above  the  mercury  in  the  trap,  the  acid  on  the  walls  of 
the  funnel  and  in  the  tube  is  washed  into  the  flask,  and  the  nitrite 
solution  to  be  analyzed  is  run  in  through  the  funnel  with  2  grm. 
of  potassium  iodide.  Sulphuric  acid  [i  13]  is  then  added  in 
amount  (5  cm.3)  sufficient  to  acidify  the  contents  of  the  flask, 
and  potassium  carbonate  is  then  added,  in  solution,  to  alkalinity 
or  until  free  iodine  has  been  taken  up.  The  mixture  is  boiled  for 
five  minutes  to  expel  nitrogen  dioxide  and  then  cooled,  and  the 
residual  arsenite  is  titrated  with  decinormal  iodine  in  presence 
of  starch.  In  making  the  various  additions  of  liquid  to  the  flask, 
care  is  of  course  taken  to  avoid  all  introduction  of  air. 

When  the  sulphuric  acid  is  added  to  the  alkaline  solution 
containing  the  arsenite,  iodide  and  nitrite,  iodine  is  set  free 
locally,  but  this  is  at  once  acted  upon  by  the  alkaline  arsenite,  so 
that  finally,  when  the  acid  reaction  is  reached,  there  is  only  a 
small  amount  still  free,  and  the  possibility  of  a  loss  of  iodine  by 
volatilization  is  reduced  to  a  minimum. 

The  table  gives  the  record  of  experiments  made  in  this  manner 
upon  a  solution  of  commercial  sodium  nitrite,  standardized  by 
treatment  with  potassium  permanganate  and  oxalic  acid  in  acid 
solution,  according  to  the  procedure  of  Kinnicut  and  Nef,f  the 

*  I.  K.  Phelps,  Am.  Jour.  Sci.,  [4],  xvii,  198. 
t  Am.  Chem.  Jour.,  v,  388. 


NITROGEN 


271 


natural  error  of  which  is  one  of  deficiency,  as  is  evidenced  by 
the  odor  of  nitrogen  oxides  observed  when  even  a  very  dilute 
solution  of  a  nitrite  is  acidified. 

Decomposition  by  Iodide  and  Ar senile  in  Acid  Solution:    Titration  of  Residual 
Ar senile  in  Alkaline  Solution. 


NaNO2  taken. 

Oxygen  value 
of  As2Os  taken. 

Oxygen  value 
of  As2O3  found. 

Error  on 
oxygen. 

Error  on 
NaNO2. 

grm. 

grm. 

grm. 

grm. 

grm. 

0.0958 

O.OI2OO 

O  .  00064 

+O.OOO25 

+O.OOII 

0.0958 

O.OI2OO 

o  .  00066 

+O.OO024 

+O.OOIO 

o.  1916 

0.03200 

o  .  00965 

+0.00017 

+0.0007 

o.  1916 

0.03200 

o  .  00965 

+0.00017 

+0.0007 

0.3832 

o  .  05600 

O.OII2O 

+0.00043 

+0.0018 

0.3832 

o  .  05600 

O.OIIlS 

+0.00045 

+0.0019 

0.6716 

o  .  08000 

O.OOl6o 

+0.00076 

+0.0033 

0.6716 

o  .  08000 

0.00158 

+0.00078 

+0.0034 

o.  1916 

0.03280 

O.OIOO3 

+0.00062 

+0.0027 

The  Estimation  of  Nitrites,  and  of  Nitrites  and  Nitrates  in  One 

Operation. 

Determination  By  the  action  of  manganous  chloride  in  hydro- 
of  Nitrites.  chloric  acid  upon  a  nitrite,  passing  the  products  of 
action  into  potassium  iodide,  and  collecting  the  residual  nitrogen 
dioxide,  Roberts*  has  been  able  to  estimate  the  nitrite  both 
from  the  iodine  set  free  and  from  the  volume  of  nitrogen  dioxide 
evolved. 

The  operation  is  conducted  in  a  slightly  modified  form  of  the 
apparatus  employed  in  the  estimation  of  nitrates. f  This  consists 
of  a  retort,  an  absorption  system  charged  with  potassium  iodide, 
a  Hempel  burette  used  for  the  collection  and  measurement  of 
the  residual  gas  over  sodium  hydroxide,!  and  a  carbon  dioxide 
generator  for  sweeping  the  gas  to  the  burette.  In  making  the 
analysis,  the  air  must  be  thoroughly  driven  out  of  the  apparatus 
before  the  nitrite  is  introduced,  as  the  carbon  dioxide,  passing 
over  the  solution,  decomposes  it.  Accordingly,  carbon  dioxide 
is  first  passed  through  the  apparatus  for  some  time,  then  the 
nitrite  is  introduced  through  the  funnel  tube  and  rinsed  in  with 

*  Charlotte  F.  Roberts,  Am.  Jour.  Sci.,  [3],  xlvi,  231. 
t  See  page  260. 
t  See  page  262. 


272 


METHODS  IN  CHEMICAL  ANALYSIS 


a  little  water,  followed  by  the  manganous  chloride  solution,  care 
being  taken  that  the  water  shall  not  exceed  one-third  of  the  total 
volume  of  the  liquid,  according  to  the  precaution  shown  to  be 
necessary  by  Gooch  and  Gruener.*  Working  in  this  way  with 
a  solution  of  specially  prepared  sodium  nitrite,  the  following 
results  were  obtained : 

Decomposition  by  Hydrochloric  Acid  and  Manganese  Chloride:    Titration  of 
Iodine  Set  Free:  Measurement  of  Nitrogen  Dioxide  Liberated. 


Volume  taken. 

NaNO2  determined 
by  KMnO4.* 

NaN02  reckoned 
from  NO. 

NaNO2  reckoned 
from  iodine. 

cm.8 

grm. 

grm. 

grm. 

10 
10 

IS 

0.0463 
0.0460 
0.0704 

0.0456 
o  .  0460 

0.0708 

0.0450 
o  .  0470 

O.O722 

15 

IS 

0.0701 
0.0688 

o  .  0704 
o  .  0696 

0.0722 
0.0695 

Determination 
of  Nitrites 
and  Nitrates. 


*  Process  of  Kinnicut  and  Nef . 

Roberts  f  has  also  shown  that  in  treating  a  mix- 
ture of  nitrite  and  nitrate  according  to  the  method 
just  described  for  the  determination  of  nitrites,  the 
measure  of  the  nitrogen  dioxide  and  the  estimation  of  liberated 
iodine  afford  data  for  the  calculation  of  the  nitrite  and  nitrate  in 
the  mixture. 

Representing  the  weight  of  nitric  oxide  found  by  a,  and  the 
weight  of  iodine  found  by  b,  and  letting  x  equal  the  amount  of 
nitric  acid  operated  upon,  and  y  the  amount  of  nitrous  acid, 

30.01  ,  30.01 

—  x  +  -—  -y  =  a, 

63.02  47.02^ 


and 
whence 


380.76       .   126.92          , 

-  x  H —  y  =  b; 

63.02  47.02  * 

x  =  0.248  b  —  1.051  a, 
y  =  2.35  a  -0.1856. 


Results  calculated  from  data  furnished  by  experiments  made 
in  the  manner  described  are  given  below. 


*  See  page  263. 
f  Loc.  cit. 


NITROGEN 


273 


Nitrite  and  Nitrate tin  One  Operation. 


NaNO2  taken, 
grm. 

NaNO2  found, 
grm. 

Error, 
grm. 

KNO3  taken, 
grm. 

KNOj  found, 
grm. 

Error, 
grm. 

0.0702 

0.0718 

+O.OOI6 

O.  IOOO 

O.IOOO 

0.0000 

0.0702 

0.0712 

-j-o.ooio 

O.  IOOO 

0.0999 

—  0.0001 

0.0702 

0.0710 

+0.0008 

O.  IOOO 

0.1004 

+0.0004 

0.0702 

0.0698 

—0.0004 

O.IOOO 

O.  IOI2 

+O.OOI2 

o  .  0468 

0.0453 

—0.0015 

O.IOOO 

0.0994 

—0.0006 

o  .  0468 

o  .  0444 

—  0.0022 

0.0500 

0.0513 

+0.0013 

In  calculating  these  results  atomic  weights  were  used,  which 
differ  somewhat  from  those  now  in  vogue,  but  the  differences 
thus  introduced  are  not  significant  where  the  inevitable  irregu- 
larities are  so  considerable. 


The  Estimation  of  Nitrates  and  Chlorates  in  One  Operation. 

A  method  for  the  determination  of  chlorates  which  has  long 
been  in  common  use  consists  in  the  treatment  of  those  compounds 
with  hydrochloric  acid,  the  passing  of  evolved  chlorine  into 
potassium  iodide,  and  the  determination  of  liberated  iodine  by 
titration  with  sodium  thiosulphate.  This  method  is  analogous 
to  the  method  proposed  by  Gooch  and  Gruener*  for  the  determi- 
nation of  nitrates,  excepting  that  in  the  latter  case  the  presence 
of  manganese  chloride  is  essential.  In  the  case  of  the  nitrate, 
however,  there  is  a  second  product,  nitrogen  dioxide,  which  may 
be  collected  and  measured,  as  in  the  process  of  treating  a  nitrate 
with  ferrous  chloride  in  hydrochloric  acid,  described  by  Roberts,  f 
By  combining  the  determination  of  the  iodine  evolved  by  the 
action  of  the  products  of  decomposition  upon  potassium  iodide 
with  the  measurement  of  nitrogen  dioxide  evolved,  Roberts  J  has 
been  able  to  effect  the  estimation  of  chlorates  and  nitrates  in  a 
single  operation  involving  distillation  with  a  solution  of  manga- 
nous  chloride  in  hydrochloric  acid,  the  amount  of  nitrogen  dioxide 
found  giving  the  amount  of  nitrate,  and  the  iodine  liberated  meas- 
uring both  nitrate  and  chlorate. 

The  operation  is  carried  out  with  a  slightly  modified  form  of 

*  See  page  263. 
f  See  page  260. 
J  Charlotte  F.  Roberts,  Am.  Jour.  Sci.,  [3],  xlvi,  231. 


274 


METHODS  IN  CHEMICAL  ANALYSIS 


the  apparatus  employed  in  the  process*  for  the  estimation  of 
nitrates  to  which  reference  has  been  made. 

In  this  apparatus,  Fig.  21,  a  small  retort  fitted  with  a  hollow 
ground-glass  stopper  prolonged  beneath  in  a  tube,  and  joined 
above  with  two  branching  tubes,  one  for  the  admission  of  carbon 
dioxide,  and  the  other,  attached  to  a  funnel  tube  with  stopcock, 
for  the  admission  of  liquids  without  introduction  of  air,  is  con- 
nected with  a  small  condenser,  which  in  turn  is  attached  to  an 
absorption  apparatus  containing  potassiuni  iodide,  and  this  with 
a  Hempel  burette  containing  a  strong  solution  of  sodium  hy- 
droxide. In  treating  the  mixture  of  chlorate  and  nitrate,  two 
Will  and  Varrentrapp  bulbs  and  generally  a  Geissler  bulb  con- 
taining potassium  iodide  are  employed  as  the  absorption  system 
to  make  sure  that  no  chlorine  shall  escape. 

The  mixture  of  chlorate  and  nitrate  is  introduced  into  the 
retort,  the  air  is  driven  out  by  carbon  dioxide,  and  then  the 
solution  of  manganous  chloride  in  hydrochloric  acid  is  added 
through  the  funnel  tube.  The  liquid  becomes  dark  at  once,  but 
a  short  heating  suffices  to  restore  it  to  its  original  clear,  light- 
green  color.  When  this  has  been  accomplished,  a  current  of 
carbon  dioxide  is  passed  through  the  apparatus,  the  bulbed  tubes 
are  disconnected,  and  their  contents  titrated  with  sodium  thio- 
sulphate.  The  volume  of  the  gas  collected  in  the  burette  is 
noted,  and  the  existing  barometric  and  thermometric  conditions, 
from  which  the  weight  of  nitrate  may  be  calculated. 

Following  are  the  results  obtained  in  tests  of  the  method. 

Nitrates  and  Chlorates  in  One  Operation. 


KC1O3  taken, 
grm. 

KC1O3  found, 
grm. 

Error.  • 
grm« 

KN03  taken, 
grm. 

KN03  found, 
grm. 

Error, 
grm. 

O.  IOOO 

0.0990 

—  O.OOIO 

O.  IOOO 

0.0995 

—0.0005 

0.0500 

o  .  0484 

—  0.0016 

0.0500 

o  .  0498 

—  O.O002 

0.0500 

o  .  0496 

—0.0004 

0.0500 

0-0515 

+0.0015 

0.0500 

o  .  0494 

—  0.0006 

0.0500 

o  .  0508 

+0.0008 

0.0500 

0.0493 

—  O.OOO7 

O.  IOOO 

0.0987 

—0.0013 

O.  IOOO 

0.0995 

—  0.0005 

O.IOOO 

0.1007 

+0.0007 

O.IOOO 

o  .  0980 

—  O.OO2O 

0.0300 

0.0305 

+0.0005 

O.IOOO 

0.0990 

—  O.OOIO 

O.IOOO 

0.1006 

+0.0006 

0.0300 

0.0293 

—  O.OOO7 

See  page  261. 


NITROGEN  275 

The  Qualitative.  Separation  and  Detection  of  Ferrocyanides,  Ferri- 
cyanides  and  Sulphocyanates. 

The  ordinary  method  of  testing  for  ferrocyanides,  ferricyanides 
and  sulphocyanates  by  means  of  ferric  and  ferrous  salts  leaves 
little  to  be  desired  in  point  of  delicacy  when  the  substances  are 
not  present  together.  When,  however,  a  sulphocyanate  and 
ferrocyanide  occur  together  the  colors  tend  to  mask  each  other, 
and  various  methods  to  obviate  the  difficulty  have  been  sug- 
gested, such  as  bleaching  the  red  ferric  sulphocyanate  by  mer- 
curic chloride,  and  distilling  the  sulphocyanic  acid  before  testing 
for  that  acid.  In  testing  for  a  ferrocyanide  in  the  presence  of  a 
ferricyanide,  the  formation  of  the  deep-blue  color  with  the  ferric 
salt  or  ferrous  salt  has  generally  been  considered  of  sufficient 
delicacy  for  all  practical  purposes. 

Browning  and  Palmer*  have  attempted  the  separation  of  these 
substances  from  one  another,  as  well  as  their  detection. 
The  Ferro-  Potassium  ferrocyanide  has  long  been  mentioned 

cyanogen  ion.  as  a  precipitant  of  the  ferrocyanides  of  the  rare  earth 
elements,  cerium,  thorium,  yttrium,  zirconium,  etc.,  while  it  is 
also  known  that  ferricyanides  of  these  elements  are  soluble. 
These  facts  suggested  the  'use  of  some  member  of  the  above- 
mentioned  group  as  a  precipitant  of  the  ferrocyanogen  ion,  and 
selection  was  made  of  a  soluble  salt  of  thorium  as  perhaps  the 
most  satisfactory  and  available.  Experience  shows  that  upon 
adding  a  few  drops  of  a  10  per  cent  solution  of  thorium  nitrate 
to  the  solution  of  a  ferrocyanide  faintly  acidified  with  acetic  acid 
it  is  possible  to  detect  by  the  cloudiness  produced  so  little  as 
I  part  of  the  ferrocyanide  in  500,000  parts  of  solution.  Alkali 
acetates  tend  to  decompose  the  thorium  ferrocyanide  into  soluble 
products,  but  the  difficulty  may  be  overcome  by  addition  of 
thorium  salt  or  hydrochloric  acid.  Neither  potassium  ferricya- 
nide nor  potassium  sulphocyanate  to  the  amount  of  o.i  grm.  in 
10  cm.3  interferes  with  this  test. 

The  Ferricyano-  In  making  the  choice  of  a  precipitant  for  the  ferri- 
genion.  cyanogen  ion  with  a  view  to  subsequent  testing  for 

the  sulphocyanogen  ion  by  the  ferric  salt,  some  reagent  giving  a 
colorless  solution  is  preferable.  Salts  of  the  elements  zinc  and 
cadmium  meet  this  condition,  and  cadmium  salts  prove  to  be  the 

*  Philip  E.  Browning  and  Howard  E.  Palmer,  Am.  Jour.  Sci.,  [4],  xxiii,  448. 


276  METHODS  IN  CHEMICAL  ANALYSIS 

more  delicate.  It  is  found  that  o.oooi  grm.  of  the  ferricyanide 
may  be  readily  detected  in  from  5  cm.3  to  10  cm.3  of  water 
acidified  with  acetic  acid  even  when  o.i  grm.  of  potassium  sul- 
phocyanate  is  present. 

Both  thorium  ferrocyanide  and  cadmium  ferricyanide  present 
difficulty  in  filtering  on  account  of  finely  divided  condition;  but 
this  difficulty  is  met  by  mixing  with  the  precipitate  fine-shredded 
asbestos  and  shaking. 

The  Ferrocyano-  ^ne  method  recommended  for  the  separation  and 
gen  ion,  the  detection  of  the  ferrocyanogen,  ferricyanogen  and 

fonT^rthfsui-8111?110^^0^11  ions  is  as  follows: 
phocyanogen  ion  I.  The  solution  to  be  tested,  preferably  dilute 
ixtures.  antj  akout  ^  cm  3  to  10  cm.3  in  volume,  is  acidified 
faintly  with  acetic  acid  or  hydrochloric  acid  and  treated  with  a 
soluble  thorium  salt  to  complete  precipitation.  To  the  liquid 
and  suspended  thorium  ferrocyanide  finely  shredded  asbestos  is 
added.  The  whole  is  agitated  and  thrown  on  a  filter,  and  the 
precipitate  is  washed  with  a  little  water.  The  washed  precipitate 
is  decomposed  by  strong  sodium  hydroxide  on  the  filter,  the  clear 
filtrate  is  acidified  with  hydrochloric  acid  and  the  test  for  the 
ferrocyanogen  ion  is  made  with  ferric  chloride. 

II.  The  filtrate  from  the  thorium  ferrocyanide  is  treated  with 
a  soluble  cadmium  salt  to  complete  precipitation  of  the  cadmium 
ferricyanide,  which,  after  the  addition  of  the  asbestos,  is  filtered 
and  washed.     The  cadmium  ferricyanide  on  the  filter  is  decom- 
posed by  sodium  or  potassium  hydroxide,  and  the  solution  is 
filtered  and  tested  with  a  ferrous  salt. 

III.  The  filtrate  from  the  cadmium  ferricyanide  is  acidified 
with  hydrochloric  acid  and  treated  with  ferric  chloride,  which 
gives  the  red  ferric  sulphocyanate. 

The  table  on  page  277  gives  results  of  practical  tests  of  this 
procedure. 

The  Gravimetric  Determination  of  Sulphocyanates. 

Van  Name*  has  shown  that  while  the  sulphocyanate  of  silver, 
unlike  that  of  copper,  is  readily  soluble  in  an  excess  of  ammonium 
or  alkali  sulphocyanates,  which  for  this  reason  may  not  be  used 
to  precipitate  silver  for  gravimetric  estimation,  the  reverse 
process,  the  precipitation  of  a  soluble  sulphocyanate  by  an  excess 
*  R.  G.  Van  Name,  Am.  Jour.  Sci.,  [4],  x,  454. 


NITROGEN 


277 


of  silver  nitrate,  furnishes  a  convenient  means  of  standardizing 
sulphocyanate  solutions  and  in  general  for  estimating  sulpho- 
cyanic  acid. 


K4FeC6N6 

present. 

grm. 

K3FeC6N6 
present. 

grm. 

KSCN 
present. 

grm. 

Indication. 

Tests  for  K4FeC6N6  only. 

O.OOIO 

0.0005 

O.OOO2 
0.0001 

O.I 
O.I 
O.  I 
O.  I 

O.  I 
O.I 
0.  I 
0.  I 

Distinct. 
Distinct. 
Distinct. 
Distinct. 

Tests  for  K3FeC6N6  only. 

O.I 
O.I 
O.I 
O.  I 

O.OOIO 
0.0005 
O.OOO2 
0  .  OOOI 

O.  I 
O.  I 
0.  I 
0.  I 

Distinct. 
Fairly  distinct. 
Faint. 
Very  faint. 

Tests  for  KSCN  only. 

O.I 
O.I 
O.  I 
O.  I 

O.I 
O.I 
O.I 
O.I 

O.OOIO 

0.0005 

O.OOO2 
O.OOOI 

Distinct. 
Distinct. 
Distinct. 
Distinct. 

Tests  for  K4FeC6N6,  K3FeC6N6  and  KSCN. 


0 
0 
0 

OIOO 

0050 

OOIO 

O.OIOO 

0.0050 

O.OOIO 

0 
0 
0 

OIOO 

0050 

OOIO 

Good 
Good 
Good 

tests  for 
tests  for 
tests  for 

K 

K 
K 

4FeC6N6, 
4FeC6N6, 
4FeC6N6, 

K3FeC6N6, 
K3FeC6N6, 
K3FeC6N6, 

KSCN. 
KSCN. 
KSCN. 

Tests  of  mixtures  unknown  to  analyst. 


O   OOIO 

O   OOIO 

Found 

K4FeCbN6, 

K3FeC6N6. 

O   OOIO 

O   OOIO 

Found 

K4FeC6N6, 

KSCN. 

O.OOIO 

O.OOIO 

O.OOIO 

Found 

K4FeC6N6, 

K3FeC6N6, 

KSCN. 

When  freshly  precipitated  the  sulphocyanate  of  silver  re- 
sembles the  chloride  in  appearance,  but  when  allowed  to  stand 
a  few  hours  becomes  finely  granular  and  is  very  easily  filtered 
and  washed.  It  may  be  safely  dried  upon  an  asbestos  filter  at 
110°  to  120°  to  a  constant  weight  corresponding  to  the  theoretical 
constitution;  but  at  a  somewhat  higher  temperature  is  decom- 
posed, leaving  a  residue  of  silver  sulphide. 


278 


METHODS  IN  CHEMICAL  ANALYSIS 


To  the  neutral  solution  of  the  sulphocyanate  in  approximately 
100  cm.3  of  water,  silver  nitrate  in  solution  is  added  in  excess. 
The  precipitate  is  collected  upon  asbestos  in  a  platinum  crucible, 
washed  with  cold  water  and  dried  to  a  constant  weight  at  115°, 
the  drying  requiring  usually  between  two  and  three  hours.  The 
filtering  is  facilitated  by  allowing  a  few  hours  for  the  precipitate 
to  settle;  but  this  is  by  no  means  essential,  as  it  is  easy  with  a 
little  care  to  obtain  a  clear  filtrate  even  when  the  filtering  is 
performed  at  once. 

In  the  following  table  are  results  obtained  by  this  procedure 
with  like  volumes  of  a  solution  of  pure  ammonium  sulphocyanate 
free  from  chloride. 

Gravimetric  Determination  of  Sulphocyanates . 
Final  Volume  of  Liquid  150  cm.3. 


NH4SCN. 
cm.3 

AgNO3. 
cm.3 

Excess  of  AgNO3. 
cm.3 

AgSCN  found, 
grm. 

25 
25 
25 
25 
25 

25 
25 
25 
25 
25 

25-3 
25-3 
25-4 
25-4 
30-4 

0.15 
0.15 
0.25 
0.25 
5-25 

0.4372 
0.4376 
0-4373 
0-4375 
0.4382 

o  .  4366 
0.4381 

0-4373 
0.4372 
0.4369 

Rough  excess. 
Rough  excess. 
Rough  excess. 
Rough  excess. 
Rough  excess. 

Mean  0.4374 

The  mean  of  the  weights  of  silver  sulphocyanate,  0.4374  grm-» 
is  equivalent  to  0.2006  grm.  of  ammonium  sulphocyanate  for 
every  25  cm.3  of  solution.  Four  titrations  of  the  same  solution 
by  Volhard's  method,  against  a  silver  nitrate  solution  whose 
standard  had  been  fixed  by  gravimetric  determination  as  silver 
chloride,  gave  as  a  mean  result  0.2003  grm.  of  ammonium  sulpho- 
cyanate for  25  cm.3  of  solution.  The  agreement  between  these 
two  values  is  within  the  possible  error  of  the  Volhard  standard. 

It  is,  therefore,  evident  that  the  standard  of  a  sulphocyanate 
solution,  free  from  chloride,  obtained  in  the  above  way  may 
safely  be  employed  for  the  estimation  of  unknown  amounts  of 
silver  by  Volhard's  method,  as  well  as  for  other  purposes. 


NITROGEN 


279 


Volumetric  Estimation  of  Sulphocyanates  by  Potassium 
Permanganate. 

When  a  solution  of  a  sulphocyanate  is  acidified  with  sulphuric 
acid  and  titrated  with  potassium  permanganate  in  the  usual 
manner,  the  end-point  is  sharp,  but  the  results,  calculated  from 
the  equation 

HSCN+3  O  +H2O  =  HCN+H2SO4, 

are  invariably  low,  the  magnitude  of  the  error  varying  greatly, 
as  the  following  tables  show,  with  the  time  occupied  in  titrating, 
temperature,  dilution,  and  amount  of  shaking.  The  principal 

Taken  for  each  experiment:  50  cm.3  of  approximately  n/6$  NH4SCN,  equiva- 
lent to  46.61  cm.3  of  the  KMnO4  solution. 


H2S04 
[1:1]. 

cm.s 

Concentra- 
tion of 
NH<SCN 
before 
titration. 

KMnO4 

used. 

cm.8 

Error, 
cm.* 

Mode  of  adding  KMnO4. 

5 

W/70 

42.75 

-3-86 

(  Slow;  little  shaking  except  at  the 
(      end. 

5 

w/70 

44-43 

-2.18 

j  Rapid;    little  shaking  except  at 
I      the  end. 

5 

w/7o 

41.41 

-5.20 

I  Very  slow  (50-100  drops  per  min- 
(      ute)  with  much  shaking. 

t  Very    rapidly    to    42    cm.3    with 

5 

w/7o 

42.78 

-3-83 

much  shaking;    last  0.78  cm.8 

(      slowly. 

(  Very  rapidly  to  44  cm.3  without 

5 

w/7o 

44.63 

—  1.98 

shaking;     last   0.63    cm.3    with 

(      shaking. 

5 

w/7o 

42.96 

-3-65 

j  Fairly    slow,    moderate    shaking 
(      throughout. 

Temperature  before  titrating  6o°-65°. 

Taken  for  each  experiment:    15  cm.3  of  approximately  w/io  KSCN,  equiva- 
lent to  87.30  cm.3  of  the  KMnO4  solution. 


H,S04 

15 

W/2O 

83.92 

-3.38 

(  Very  rapid;   little  shaking  except 
I      at  the  end. 

15 

W/20 

82.20 

-5.10 

Slow,  much  shaking. 

30 

w/30 

84.82 

-2.48 

(  Very  rapid;   little  shaking  except 
(      at  the  end. 

!    Medium  rate  (averaging  about 

30 

w/30 

83.02 

-4.28 

15  cm.3  per  minute  until  near  the 

end),  with  shaking. 

280  METHODS  IN  CHEMICAL  ANALYSIS 

cause  of  the  deficiencies  is  probably  the  fact  noted  by  Klason  * 
and  others,  that  sulphocyanic  acid  undergoes  decomposition  in 
water  solution  in  the  presence  of  an  inorganic  acid. 

It  has,  however,  been  shown  by  Van  Name  f  that  if  the  neutral 
solution  of  the  sulphocyanate  be  slowly  added  to  a  measured  and 
sufficiently  large  excess  of  permanganate,  previously  acidified 
with  sulphuric  acid,  and  the  excess  of  permanganate  then  esti- 
mated with  oxalic  acid,  the  results  are  close  to  the  theory. 
Under  these  conditions  the  sulphocyanic  acid  is  not  liberated 
until  in  actual  contact  with  the  permanganate,  and  side  reactions 
are  thus  avoided.  -To  insure  this  result  there  must  at  all  times 
be  a  sufficient  excess  of  permanganate  present  to  give  the  solu- 
tion a  deep-red  color.  Moreover,  since  manganous  salts  reduce 
permanganate,t  no  considerable  amount  of  manganous  salt  can 
be  produced  by  the  reaction  with  the  sulphocyanate  until  after 
all  the  permanganate  has  been  converted  into  the  brown  oxide. 
The  excess  of  permanganate  originally  taken  must  therefore 
be  at  least  five-thirds  of  the  quantity  actually  required  for  the 
oxidation  of  the  sulphocyanate,  in  order  to  insure  the  presence 
of  unchanged  permanganate  throughout.  In  practice  a  larger 
excess  is  desirable,  at  least  twice  —  or  better  three  times  —  the 
theory,  and  amounts  as  large  as  ten  times  the  theory  may  safely 
be  used. 

It  is  important,  however,  to  remove  the  excess  of  permanga- 
nate without  delay,  since  the  instantaneous  oxidation  of  the  sul- 
phocyanate is  followed  by  a  slow  loss  of  permanganate,  probably 
due  to  an  oxidation  beyond  the  cyanide  stage,  the  rate  of  loss 
being  greatest  when  the  solution  is  concentrated  and  the  excess 
of  permanganate  large.  This  error  can  be  avoided  by  having 
the  oxalate  solution  in  readiness  for  rapid  addition  immediately 
after  the  sulphocyanate,  or  by  at  once  precipitating  the  perman- 
ganate by  adding  an  excess  of  a  manganous  salt,  thus  removing 
the  necessity  for  further  haste. 

The  following  procedure  is  recommended :  A  volume  of  stand- 
ard n/io  permanganate  solution,  sufficient  to  oxidize  at  least 
two  and  one-half  times  the  quantity  of  sulphocyanate  actually 

*  Jour,  prakt.  Chem.  (N.  F).,  xxxvi,  74. 

t  R.  G.  Van  Name,  Dissertation,  Yale  University,  1902.  The  details  of  the 
procedure  as  here  given  and  the  three  tables  are  taken  from  an  unpublished 
portion  of  this  thesis,  on  file  since  1902  in  the  Yale  University  Library. 

{  The  so-called  Guyard  reaction,  Mn2O7  -f-  3  MnO  =  5  MnO2. 


NITROGEN 


28l 


to  be  determined,  is  measured  out  and  acidified  with  about  one- 
tenth  of  its  volume  of  [1:1]  sulphuric  acid.  To  this  the  sul- 
phocyanate  in  neutral  or  faintly  alkaline  solution,  one-fiftieth 
normal  or  .stronger,*  is  added  rapidly,  a  few  drops  at  a  time,  with 
thorough  stirring,  and  is  immediately  followed  by  concentrated 
manganous  sulphate  solution  in  quantity  sufficient  to  precipitate 
all  the  residual  permanganate.  The  manganese  oxides  are  then 
dissolved  (with  or  without  previous  warming  of  the  liquid  to 
75°)  by  a  measured  amount  of  standard  ammonium  oxalate  or 
oxalic  acid,  and  the  titration  to  color  completed  with  the  perman- 
ganate. It  is  best  to  leave  for  this  purpose  a  few  cubic  centi- 
meters of  permanganate  in  the  burette  when  the  first  portion  is 
taken,  thus  avoiding  unnecessary  burette  readings. 

The  following  analyses  were  carried  out  by  the  method  just 
described,  the  ratio  of  permanganate  taken  to  that  actually 
required  being  about  ten  to  one  in  the  first  two  experiments,  and 
slightly  above  two  to  one  in  the  last  four. 

Determination  of  Sulphocyanates  by  Potassium  Permanganate. 


KMnO4. 

H2S04[i:i]. 

HSCN  equiva- 
lent to  KSCN 

KMnO4  used. 

HSCN  found. 

Error. 

taken. 

cm.» 

cm.* 

grm. 

cm.8 

grm. 

grm. 

45 

5 

0.00482 

4-H 

O  .  00486 

+0  .  00004 

45 

5 

0.00482 

4.10 

0.00485 

-{-0.00003 

45 

5 

0.01345 

12.  2O 

0.01343 

—  O.OOOO2 

45 

5 

0.02408 

20.  l6 

0.02385 

—  0.00023 

45 

5 

0.02408 

2O.  21 

0.02391 

—0.00017 

95 

10 

0.04816 

40.50 

0.04791 

—  0.00025 

195 

20 

0.09631 

81.11 

0-09595 

—0.00036 

When  the  amount  of  sulphocyanic  acid  to  be  estimated  is 
entirely  unknown,  the  color  of  the  solution  must  be  carefully 
watched,  and  if  the  red  color  of  the  permanganate  grows  weak 
the  addition  of  the  sulphocyanate  must  be  stopped,  a  further 
quantity  of  permanganate  added,  and  the  process  completed  in 
the  usual  way. 

On  account  of  the  large  volume  of  permanganate  required,  the 
method  is  best  adapted  for  the  estimation  of  small  amounts  of 
sulphocyanic  acid. 

*  Higher  dilutions  may  be  employed,  but  the  excess  of  permanganate 
should  be  proportionately  increased. 


282 


METHODS  IN  CHEMICAL  ANALYSIS 


PHOSPHORUS. 

The    Determination    of  Phosphoric   Acid    by    Precipitation    as 
Ammonium  Magnesium  Phosphate  and  Weighing  as 
Magnesium  Pyrophosphate. 

Gooch  and  Austin*  have  shown  that  the  precipitation  of  a 
soluble  phosphate  by  the  magnesia  mixture  is  practically  com- 
plete in  faintly  ammonical  solutions  even  when  very  dilute  and 
charged  with  large  amounts  of  ammonium  chloride,  provided 
the  magnesia  mixture  is  present  in  sufficiently  large  excess; 
also  that  ammonium  salts  tend  to  produce  an  ammonium  mag- 
nesium phosphate  richer  in  ammonia  and  phosphoric  acid  and 
poorer  in  magnesium  than  the  normal  salt,  NH4MgPO4,  while 
an  excess  of  the  magnesia  mixture  tends  to  produce  excess 
of  magnesium  in  the  precipitated  phosphate.  The  results  of 
experiment,  recorded  below,  go  to  show  that  good  results  may 
be  expected  when  the  solution  of  the  phosphate,  containing  a 
moderate  excess  of  the  magnesium  salt  and  not  more  than  5  to 
10  per  cent  of  ammonium  chloride,  is  precipitated  by  making 
it  slightly  ammoniacal,  the  precipitate  being  washed  in  slightly 
ammoniacal  wash-water.  In  general,  however,  and  especially 
when  more  ammonium  chloride  than  this  proportion,  or  more 
magnesium  salt  than  twice  the  amount  theoretically  necessary, 

Estimation  as  Magnesium  Pyrophosphate:  Effect  of  Ammonium  Salts. 


Mg2P207 

correspond- 
ing to 
HNa2P04 
taken. 

grm. 

Mg2P207 
found. 

grm. 

Error  in 
terms  of 
Mg2P207- 

grm. 

Error  in 
terms  of 
P. 

grm. 

Volume. 
cms. 

NH4C1 
in  mag- 
nesia 
mixture. 

grm. 

NH4C1 
added. 

MgCl2. 
6H2Oin 
mag- 
nesia 
mixture. 

grm. 

I. 

grm. 

II. 

grm. 

Single  precipitation. 


0.8615 

0.8613 

—  0.0002 

—  O.OOOO5 

150 

1.68 

3-3 

0.8615 

0.8615 

O.OOOO 

0  .  OOOOO 

2OO 

1.68 

20 

3-3 

0.8615 

0.8602 

—  0.0013 

—  0.00036 

2OO 

1.68 

20 

3-3 

0.8615 

0.8561 

—  0.0054 

—0.00151 

300 

1.68 

60 

3-3 

Double  precipitation. 


0.8111 

0.8114 

+0.0003 

-|-o  .  00008 

150  100 

1.68 

3-3 

0.8615 

0.8613 

—  0.0002 

—0.00006 

150  ioo 

1.68 

3-3 

0.8615 

0.8578 

-0.0037 

—0.00103 

2OO  IOO 

1.68 

20 

3-3 

0.8615 

0.8487 

—  0.0128 

—0.00358 

2OO  IOO 

1.68 

60 

3-3 

*  F.  A.  Gooch  and  Martha  Austin,  Am.  Jour.  Sci.,  [4],  vii,  187. 


PHOSPHORUS  283 

is  present,  it  is  safer  to  decant  the  supernatant  liquid  from  the 
precipitate  (through  the  filter  to  be  used  subsequently  to  hold 
the  phosphate) ,  to  dissolve  the  precipitate  in  a  little  hydrochloric 
acid  and  reprecipitate  by  dilute  ammonia,  washing  with  faintly 
ammoniacal  wash-water. 

The  lodometric  Determination  of  Phosphorus  in  Iron. 

The  very  careful  work  of  Blair  and  Whitfield  *  shows  that  the 
ammonium  phosphomolybdate,  precipitated  under  the  condi- 
tions ordinarily  prescribed  for  the  determination  of  phospho- 
rus in  iron  or  iron  ores,  is  of  the  definite  constitution  expressed 
by  the  symbol  24  MoO3,  P2O5,  3  (NH4)2O,  2  H2O,  and  contains 
1.794  parts  of  phosphorus  to  every  100  parts  of  molybdic  an- 
hydride. It  has  been  demonstrated  in  the  work  of  Fair  banks  f 
that  phosphorus  may  be  determined  with  advantage  by  finding 
iodometrically  the  combined  molybdic  anhydride.  Molybdic  acid 
may  be  reduced  by  hydriodic  acid  to  the  condition  of  oxidation 
represented  by  the  symbol  Mo2Os  in  acid  solution, 

2  MoO3  +  2  HI  =  Mo2O5+I2+H2O; 

while  in  an  alkaline  solution  the  reduced  product  is  reoxidized 
by  standard  iodine. 

Mo2O5+I2+H2O  =  2  MoO3+2  HI. 

The  amount  of  iodine  necessary  to  reoxidize  reduced  molybdic  acid 
is  large,  and  the  amount  of  molybdic  acid  compared  with  the 
phosphorus  contained  in  the  phosphomolybdate  is  also  large,  so 
that  a  method  of  great  theoretical  accuracy  should  result  from 
the  utilization  of  these  reactions  for  the  determination  of  phos- 
phorus in  iron. 

According  to  the  procedure  described,  the  solution,  hot  less, 
than  150  cm.3  nor  more  than  300  cm.3  in  volume,  and  containing 
iron  as  the  nitrate  and  phosphorus  as  phosphoric  acid  in  presence 
of  not  too  much  free  nitric  acid,  is  heated  to  85°  and  shaken 
for  five  minutes  with  40  cm.3  of  filtered  ammonium  molybdate 
solution, t  and  filtered  on  asbestos  in  the  perforated  crucible. 

*  Jour.  Am.  Chem.  Soc.,  xvii,  747. 

t  Charlotte  Fairbanks,  Am.  Jour.  Sci.,  [4],  ii,  181. 

t  Blair  and  Whitfield,  loc.  cit.  This  solution  is  made  of  100  grm.  of  MoO3^ 
400  cm.3  of  water;  80  cm.3  of  concentrated  NH4OH  added  to  300' cm.3  of 
HNOs  (sp.  gr.  1.42),  diluted  with  700  cm.3  of  water.  > 


284  METHODS  IN  CHEMICAL  ANALYSIS 

The  precipitate  is  washed  with  10  per  cent  nitric  acid  and  then 
with  I  per  cent  potassium  nitrate.  The  asbestos  felt  is  trans- 
ferred to  a  i5O-cm.3  flask  or  narrow-based  Erlenmeyer.  The 
precipitation  flask  and  cork  are  thoroughly  washed  with  a  mix- 
ture of  5  cm.3  of  ammonia  and  10  cm.3  of  water,  and  the  wash- 
ings are  allowed  to  rinse  the  sides  of  the  perforated  crucible  — 
standing  on  a  small  funnel  —  and  so  to  run  into  the  flask.  Strong 
hydrochloric  acid  is  added,  25  cm.3,  and,  when  the  phosphorus 
does  not  exceed  0.0060  grm.,  0.5  grm.  of  potassium  iodide;  but 
when  more  phosphorus  is  present  a  little  more  potassium  iodide 
is  needed.  Experience  has  shown  that  the  iodide  present 
should  not  exceed  the  amount  theoretically  necessary  by  more 
than  a  half-gram.  The  flask  is  trapped  loosely  with  a  short 
bulbed  tube  hung  in  the  neck.* 

The  liquid  is  boiled  down  from  a  total  volume  of  40  cm.3  to 
just  25  cm.3,  easily  marked  by  two  strips  of  paper  pasted  on 
opposite  sides  of  the  flask.  If  the  solution  is  boiled  farther,  the 
molybdic  acid  is  likely  to  be  reduced  beyond  the  degree  of  oxi- 
dation indicated  by  the  symbol  Mo2O5. 

The  residue  is  cooled  and  transferred  to  a  stoppered  bottle 
(shown  in  Fig.  2)  fitted  with  a  separatory  funnel  and  a  trap  filled 
with  a  solution  of  potassium  iodide.  Through  the  stoppered  fun- 
nel- are  added  in  solution  I  grm.  of  tartaric  acid,  enough  sodium 
hydroxide  to  nearly  neutralize  the  free  acid,  followed  by  acid 
sodium  carbonate  to  complete  the  neutralization,  and  a  measured 
amount  of  iodine  in  excess  of  that  required  for  the  oxidation. 

After  neutralization  the  iodine  color  in  the  solution  should 
perceptibly  fade  within  fifteen  minutes ;  but  for  complete  oxida- 
tion the  bottle  should  be  set  aside,  out  of  sunlight,  for  an  hour  and 
a  half,  and  then  the  excess  of  the  iodine  is  titrated  with  a  stand- 
ard solution  of  arsenious  acid. 

Since  there  is  a  slight  tendency  on  the  part  of  the  iodine  to 
form  a  little  iodate  during  the  long  digestion,  it  is  wise  to  acidulate 
the  solution  in  each  case  slightly  with  dilute  hydrochloric  acid 
after  the  titration  with  the  arsenic  solution,  and  then  to  deter- 
mine by  sodium  thiosulphate  the  trace  of  iodine  which  has  taken 
the  form  of  iodate. 

In  the  following  table  of  test  experiments  the  absolute  errors 
in  terms  of  phosphorus  are  given;    and  the  percentage  errors, 
*  See.  Fig.  6,  page  6. 


PHOSPHORUS 


285 


between  the  phosphorus  taken  and  the  phosphorus  found,  re- 
ferred to  10  grm.  of  material  (the  maximum  amount  of  high- 
grade  iron  or  steel  usually  taken  for  analysis),  are  also  added. 

Reduction  of  Phosphomolybdate  by  Hydriodic  Acid:    lodometic  Determination 
of  Reduced  Molybdic  Acid. 


Amount  of  P 
taken, 
gnn. 

Amount  of  P 
found, 
grm. 

Error  on  P. 
grm. 

Error  of  P. 
Per  cent, 
grm. 

Neutralized  by 

0.002727 
o.  001812 
o  .  000909 
0.003508 
0.005454 
o.  001818 

0.003636 

o  .  000909 
0.000363 
0.008180 

0.002778 
0.001743 
0.000914 
0.003262 
0.005417 

o.  001861 

0.003716 

o  .  000988 

0.000289 
0.008179 

+0.000051 
—  0.000069 
+0.000005 
—  0.000246 
—  0.000037 
+o  .  000043 
+0.000080 
+0.000079 
—  0.000074 
—  O.OOOOOI 

+O.OOO5 
—  0.0007 

+o  .  00005 

—  O.OO2* 
—  0.0003 
+0.0004 

+o  .  0008 

+O.OOO8 
—  O.OOO7 
—  O.OOOOI 

NaHCO3. 
NaOH+NaHCO3. 
NaOH+NaHC03. 
NaOH+NaHCO3. 
NaHCO3. 
NaOH+NaHCO3. 
NaHCO3. 
NaHC03. 
NaOH+NaHCO3. 
NaHCO3. 

Obviously  accidental. 


The  Estimation  of  Phosphoric  Acid  and  Phosphorus  Precipitated 
as  Ammonium  Phosphomolybdate. 

The  method  studied  by  Randall  *  for  the  estimation  of  molyb- 
dic  acid  by  the  aid  of  the  zinc  reductor,  the  receiving  flask  charged 
with  ferric  alum,  and  the  permanganate  titration,  has  been  ap- 
plied by  himf  to  the  estimation  of  phosphorus  in  iron,  precipi- 
tated in  the  form  of  ammonium  phosphomolybdate.  According 
to  Randall's  procedure,  the  phosphomolybdate,  precipitated  in 
a  flask  and  shaken  in  the  usual  manner,  is  allowed  to  settle,  then 
filtered  on  asbestos  in  a  perforated  crucible,  and  washed  with  a 
solution  of  ammonium  acid  sulphate  (15  cm.3  ammonia,  25  cm.3 
sulphuric  acid,  I  liter  water).  The  flask  is  washed  out  with  a 
solution  of  20  cm.3  of  water  and  5  cm.3  of  ammonia,  and  this 
is  poured  on  the  asbestos  in  the  crucible.  The  molybdenum 
solution  is  acidified  with  10  cm.3  of  strong  sulphuric  acid  and 
passed  through  the  reductor  into  the  ferric  alum  solution,  pre- 
ceded by  100  cm.3  of  hot  water  and  followed  by  200  cm.3  of  the 
hot  dilute  acid  with  100  cm.3  of  water,  the  reduced  solution  being 
titrated  immediately  with  approximately  tenth  normal  permanga- 
nate. The  results  are  calculated  on  the  assumption  that  the  am- 

*  See  page  424. 

f  D.  L.  Randall,  Am.  iour.  Sci.,  [4],  xxiv,  315. 


286 


METHODS  IN  CHEMICAL  ANALYSIS 


monium  phosphomolybdate  contains  phosphorus  and  molybde- 
num in  the  proportion  given  by  the  symbol  (NH4)3i2MoO3PO4, 
and  that  the  reduction  proceeds  to  the  condition  represented 
by  the  symbol  Mo2O3.  In  the  following  table  are  shown  results 
obtained  by  this  procedure  applied  to  pure  ferric  nitrate  and  a 
known  amount  of  microcosmic  salt. 


Determination  of  Phosphorus  by  Titration  of 
Reduced  Phosphomolybdate. 


P  taken, 
grm. 

P  found, 
grm. 

Error, 
grm. 

0  .  003645 

0.003673 

+0.000028 

0.003645 

0.003697 

+0.000052 

0.003645 

0.003638 

—  0.000007 

0  .  003645 

0.003726 

+O.OOOo8l 

0.003645 

o  .  003630 

—  0.000015 

0.003645 

0.003661 

+O.OOOOI6 

The  Determination  of  Phosphoric  Acid  by  Precipitation  as  Uranyl 
Phosphate  and  Estimation  of  the  Uranium  Volumetrically. 

The  method  of  estimating  uranium  by  reduction  in  the  zinc 
redactor  and  oxidation  with  permanganate*  has  been  applied 
by  Pulmanf  to  the  determination  of  uranic  oxide,  and  so  of 
phosphoric  acid,  precipitated  as  ammonium  uranyl  phosphate. 
This  precipitate  is  filtered  with  extreme  difficulty,  but  with  an 
asbestos  felt  coated  with  the  finer  floating  particles  of  the  par- 
tially settled  emulsion  of  prepared  asbestos, it  is  possible  to 
obtain  filtrates  from  the  ammonium  uranyl  phosphate  which  are 
perfectly  clear,  though  the  process  of  filtering  and  washing  is 
slow  on  account  of  the  compactness  of  the  felt  surface  and  the 
gelatinous  nature  of  the  precipitate. 

The  process  as  worked  out  for  the  determination  of  the  phos- 
phoric acid  is  as  follows:  A  measured  amount  of  a  standard 
phosphate  solution  (containing  about  4.7  grm.  of  microcosmic 
salt  per  liter)  is  drawn  into  a  beaker,  and  a  solution  containing 
12  grm.  of  ammonium  acetate,  formed  by  neutralizing  about 
10  cm.3  of  ammonium  hydroxide  (0.90  sp.  gr.)  with  acetic  acid 


*  See  page  430. 

t  O.  S.  Pulman,  Jr.,  Am.  Jour.  Sci.,  [4],  xvi,  229. 


PHOSPHORUS 


287 


(50  per  cent),  and  from  2  cm.3  to  4  cm.3  of  free  acetic  acid  is 
added.  The  total  volume  is  made  up  to  about  150  cm.3  and 
the  solution  heated  nearly  to  boiling.  The  ammonium  uranyl 
phosphate  is  then  precipitated  by  slowly  adding  an  excess  of 
uranium  nitrate,  with  stirring,  and  the  mixture  is  boiled  gently 
for  about  twenty  minutes,  allowed  to  settle,  and  filtered  on  a 
tight  felt  of  asbestos.  The  precipitating  beaker  and  the  pre- 
cipitate are  washed  thoroughly  with  a  dilute  solution  of  ammo- 
nium acetate  containing  a  little  free  acetic  acid  (to  overcome  the 
tendency  of  the  precipitate  to  pass  through  the  filter),  and  the 
crucible  containing  the  precipitate  is  placed  in  a  glass  funnel. 
Enough  dilute  sulphuric  acid  [1:5]  is  then  added  to  dissolve 
the  precipitate  and  thoroughly  wash  out  all  the  soluble  uranium 
salt  from  the  asbestos,  the  solution  being  caught  below  as  it 
passes  through  the  crucible  and  funnel  in  the  beaker  used  for 
the  precipitation.  The  solution  is  made  up  to  a  volume  of  from 
100  cm.3  to  150  cm.3  with  dilute  sulphuric  acid  [1:5],  heated 
to  boiling.  A  few  cubic  centimeters  of  warm  dilute  sulphuric 
acid  [1:5]  are  passed  through  the  reductor  and  followed  by  the 
uranium  solution,  a  few  cubic  centimeters  more  of  the  dilute 
sulphuric  acid,  and  250  cm.3  of  hot  water.  The  contents  of  the 
flask  are  then  poured  into  a  porcelain  dish,  diluted  with  200  cm.3 
of  hot  water,  and  titrated  with  a  n/io  solution  of  potassium 
permanganate. 

Results  obtained  by  this  method  are  shown  in  the  following 
table: 

Reduction  of  Uranyl  Phosphate  and  Titration  with  Permanganate. 


P*06 

taken. 

UO3  corre- 
sponding to 
P2O5  taken. 

H2SO4 

(1.84). 

Dilution 
at  reduc- 
tion. 

KMnO4. 

UO3  found. 

Error  on 
U03. 

Error  on 
P206. 

grin. 

grm. 

cm.8 

cm.8 

cm.3 

grm. 

grm. 

grin* 

o  .  0404 

o.  1630 

25 

ISO 

II  .06 

0.1632 

+O.OOQ2 

+0.00005 

o  .  0404 

o.  1630 

25 

150 

11.03 

0.1628 

—  0.0002 

—0.00005 

O.O226 

O.O9I2 

20 

1  2O 

6.14 

o  .  0906 

—  O.OOO6 

—0.00015 

O.O226 

O.O9I2 

20 

1  2O 

6.17 

0.0911 

•^•o.oooi 

—  O.OOOO2 

0.0719 

o.  2902 

25 

ISO 

19.62 

0.2896 

—  O.OOO6 

—  0.00015 

0.0719 

0.2902 

25 

ISO 

19.61 

0.2894 

—  0.0008 

—  O.OOO2O 

The  process  for  the  determination  of  uranium  by  the  reductor 
depends  upon  the  fact  that  any  reduction  of  uranium  lower  than 
uranous  oxide  —  and  such  reduction  undoubtedly  takes  place 


288  METHODS  IN  CHEMICAL  ANALYSIS 

in  the  reductor  —  is  corrected  by  exposure  to  the  air,  the  lower 
oxide  being  rapidly  oxidized  to  exactly  the  uranous  state,  while 
the  uranous  salts  are  stable  enough  to  be  estimated  before  they 
are  oxidized  appreciably  by  atmospheric  action. 

ARSENIC,  ANTIMONY  AND  TIN. 

The  Determination  of  Arsenic  by  Precipitation  as  Ammonium 

Magnesium  Ar senate  and    Weighing  as   Magnesium 

Pyroarsenate. 

The  striking  analogy  between  the  phosphates  and  the  arse- 
nates  led  Levol*  to  undertake  the  separation  of  an  ammonium 
arsenate  corresponding  to  the  ammonium  magnesium  phos- 
phate, the  composition  of  which  Berzelius  had  given.  Levol 
states  that  ammonium  magnesium  arsenate  of  the  composition 
NH4MgAsO4.ioH2O  is  obtained  by  adding  a  solution  of  a  double 
ammonium  magnesium  salt  to  arsenic  acid,  that  it  is  a  salt 
possessing  about  the  same  degree  of  solubility  in  water,  in  am- 
moniacal  water,  and  in  ammoniacal  water  containing  magnesium 
salt,  as  the  corresponding  phosphate,  and  that  at  red  heat,  after 
carefully  drying,  it  yields  magnesium  pyroarsenate.  Several 
sources  of  error  in  this  process  have  been  pointed  out  by  many 
investigators.  First,  the  low  indications  obtained  when  the 
pyroarsenate  is  weighed  suggest  a  loss  of  arsenic  during  ignition, 
in  consequence  of  the  reducing  action  of  ammonia  evolved  in  the 
process. f  Another  possible  source  of  error  is  the  solubility  of 
ammonium  magnesium  arsenate  in  ammoniacal  wrater  and  solu- 
tions of  ammonium  salts.J  There  is  also  the  possibility  that  the 
constitution  of  ammonium  magnesium  arsenate  as  precipitated 
may  not  be  ideal,  in  consequence  of  the  action  of  ammonium 
salts  known  to  be  influential  in  determining  the  constitution  of 
the  analogous  ammonium  phosphates  of  magnesium  and  other 
elements.§ 

*  Ann.  Chim.,  [3],  xvii,  501. 

t  Wach  and  Rose,  Schweigger,  Jour.  Ch.  Phys.,  lix,  297.  Reichel,  Ann. 
Phys.,  Ixxvi,  20.  Rammelsberg,  Ber.  Dtsch.  chem.  Ges.,  xiv,  279.  Kaiser^ 
Zeit.  anal.  Chem.,  xiv,  250. 

t  Rose,  Zeit.  anal.  Chem.,  iii,  206.  Wood,  Am.  Jour.  Sci.,  (3],  vi,  368. 
Brauner,  Zeit.  anal.  Chem.,  xvi,  57. 

§  Neubauer,  Zeit.  anorg.  Chem.,  ii,  45;  Zeit.  angew.  Chem.,  1896,  435; 
Jour.  Am.  Chem.  Soc.,  xvi,  289.  Gooch  and  Austin,  Am.  Jour.  Sci.,  [4],  vi, 
233;  vii,  187;  viii,  206. 


ARSENIC,   ANTIMONY  AND   TIN  289 

The  conditions  to  be  observed  in  applying  the  method  to  the 
determination  of  arsenic  have,  therefore,  been  carefully  studied 
by  Austin.* 

It  is  shown  in  the  first  place  by  special  tests  that  the  presence 
of  ammonium  chloride  does  produce  solubility  of  the  precipitate 
thrown  down  by  magnesia  mixture,!  but  that  this  solvent  effect 
may  be  overcome  even  when  the  ammonium  chloride  present 
amounts  to  as  much  as  60  grm.  in  300  cm.3  of  total  volume, 
by  a  sufficiency  of  the  magnesia  mixture;  and,  further,  that  the 
ammonium  magnesium  arsenate  once  precipitated  may  be  safely 
washed  with  small  amounts  (25  cm.3  to  50  cm.3)  of  faintly  ammo- 
niacal  water. 

It  appears  also  that  the  ammonium  salt  induces  the  formation 
of  an  ammonium  magnesium  arsenate  too  rich  in  ammonia  to 
give  the  pyroarsenate,  Mg2As2O7,  on  ignition,  even  when  the 
precipitation  is  complete.  It  is  found,  however,  that  a  suitable 
increase  in  the  amount  of  magnesia  mixture  present  at  precipi- 
tation may  bring  about  the  formation  of  an  ammonium  magne- 
sium arsenate  of  ideal  constitution,  even  in  presence  of  a  consid- 
erable amount  of  the  ammonium  salt. 

According  to  the  most  favorable  procedure,  the  slightly  acid 
solution  of  the  arsenate  containing  no  ammonium  salts  is 
added  drop  by  drop  to  the  distinctly  ammoniacal  magnesia 
mixture,  and  a  little  more  ammonia  is  added.  The  precipitate 
is  filtered  off  as  soon  as  it  subsides,  on  asbestos  in  the  per- 
forated crucible  and  with  use  of  the  filtrate  to  effect  the  trans- 
fer, and  washed  with  about  25  cm.3  of  faintly  ammoniacal  water 
applied  in  small  portions.  After  careful  drying,  the  residue  is 
ignited  with  caution  and  weighed  as  magnesium  pyroarsenate, 
Mg2As2O7. 

When  no  ammonium  salts  are  present  an  excess  of  about 
30  cm.3  of  magnesia  mixture  in  a  total  volume  of  200  cm.3  is 
sufficient  to  form  the  arsenate  in  ideal  condition.  If  ammonium 
salts  are  present  and  conditions  prevent  their  removal,  the 

*  Martha  Austin,  Am.  Jour.  Sci.,  [4],  ix,  55. 

t  The  magnesia  mixture  is  prepared  by  dissolving  no  grm.  of  the  crys- 
tallized magnesium  chloride  in  a  small  volume  of  water,  filtering,  and  adding 
to  it  58  grm.  of  ammonium  chloride  (purified  in  solution  by  adding  bromine 
water  and  bleaching  with  ammonia),  filtering,  diluting  to  a  volume  of  2  liters, 
and  adding  enough  ammonia  — 10  c.c.  —  to  make  the  solution  smell  distinctly 
of  ammonia. 


2QO 


METHODS  IN  CHEMICAL  ANALYSIS 


slightly  acidulated  arsenate  should  be  added  gradually  to  a  very 
large  excess  (150  cm.3)  of  the  magnesia  mixture. 

Results  obtained  by  the  procedure  are  given  in  this  table. 

Determination  of  Arsenic  as  Magnesium  Pyr 'oar senate. 


Mg2As2O7  corresponding  to  As2O5. 

Magnesia 
mixture. 

cm.* 

NH4C1. 

grm. 

Taken, 
grin. 

Found, 
grm. 

Error, 
grm. 

In  absence  of  ammonium  salts. 


o  .  7843 

0.7830 

—0.0013 

50 

0.7843 

o  .  7849 

—0.0006 

50 

0.7843 

0.7841 

—  O.OOO2 

50 

. 

0.7843 

0.7843 

0  .  0000 

50 

In  presence  of  ammonium  salts. 


0.7843 

0.7763 

—0.0080 

75 

IO 

0.7843 

0.7762 

—0.0081 

75 

IO 

0.7843 

0.7832 

—  O.OOII 

IOO 

IO 

o  .  7843 

0.7838 

—0.0005 

100 

10 

0.7843 

0.7784 

-0.0059 

IOO 

20 

0.7843 

0.7810 

-0.0033 

IOO 

20 

0.7843 

o  .  7849 

-f-o  .  0006 

150 

60 

0.7843 

o  .  7846 

+o  .  0003 

150 

60 

In  no  one  of  the  many  precipitates  tested  by  silver  nitrate 
for  included  chlorides  was  more  than  a  trace  found. 

Precipitation  of  Small  Amounts  of  the  Arsenate. 


Mg2As207. 

Magnesia 
mixture. 

NH4C1. 

Taken. 

Found. 

Error. 

cm.8 

grm. 

grm. 

grm. 

grm. 

25 

0.0015 

0.0014 

—  O.OOOI 

25 

0.0015 

0.0017 

+O.OOO2 

25 

. 

0.0077 

0.0078 

+O.OOOI 

50 

0.0077 

0.0076 

—  O.OOOI 

25 

IO 

0.0015 

0.0013 

—  O.OOO2 

25 

10 

0.0077 

0.0070 

—0.0007 

To  effect  the  precipitation  of  amounts  of  arsenate  so  small  as 
not  to  be  at  once  precipitable,  recourse  may  be  taken  to  a  proc- 


ARSENIC,  ANTIMONY  AND  TIN  291 

ess  of  freezing  and  melting,  in  which  the  magnesium  arsenate 
is  made  insoluble  in  the  freezing  and  continues  to  be  insoluble 
when  the  medium  melts.  This  is  best  effected  by  putting  the 
solution  in  a  platinum  dish,  surrounding  the  dish  with  a  mix- 
ture of  ice  and  salt  until  the  mass  is  solid,  and  then  allowing  the 
mass  to  melt  at  the  room  temperature.  Results  obtained  in  this 
manner  by  Gooch  and  Phelps*  are  given  in  the  table. 

The  lodometric  Estimation  of  Arsenic  Acid. 

Holthoff's  development  of  Mohr's  suggestion  relative  to  the 
reduction  of  arsenic  acid  to  the  lower  condition  of  oxidation  by 
the  action  of  sulphurous  acid,f  with  the  demonstration  that 
arsenic  acid  can  be  evaporated  even  to  dryness  in  presence  of 
hydrochloric  acid  without  danger  of  significant  volatilization,  has 
placed  the  analysis  of  ordinary  compounds  of  arsenic  within  the 
scope  of  Mohr's  classical  and  exact  method  of  determination  by 
titration  with  iodine.  As  Holthoff  left  the  method,  it  is  satis- 
factory so  far  as  regards  accuracy,  and  as  modified  by  McCay,{ 
who  substitutes  for  the  four  hours'  digestion  heating  for  one  hour 
in  a  pressure  bottle,  is  eminently  successful.  Gooch  and  Brown- 
ing! have  still  further  shortened  the  process  of  reduction  of 
arsenic  acid  by  making  use  of  hydriodic  acid  as  the  active  agent 
instead  of  sulphurous  acid. 

A   method    for   the   determination   of   iodine   in 

Reduction  by        ,,.,,, 

haloid  salts  based  upon  the  action  of  arsenic  acid, 


and  Oxidation     m  ^e  presence  of  sulphuric  acid,  according  to  the 

by  Iodine  in 

Alkaline  Solu-      equation, 

tion'  H3As04  +  2  HI  =  HsAsOs  +  H2O  +  I2, 

the  iodine  being  completely  volatilized,  but  leaving  behind  in 
the  arsenious  acid  produced  by  the  action  the  record  of  the 
amount  of  hydriodic  acid  originally  present,  is  described  else- 
where. ||  This  reaction  is  in  the  present  case  utilized  conversely, 
and  potassium  iodide  in  excess,  in  presence  of  sulphuric  acid,  is 
employed  to  bring  about  the  reduction  of  the  arsenic  acid  to 
arsenious  acid,  which  may  be  determined,  after  neutralization,  by 

*  F.  A.  Gooch  and  M.  A.  Phelps,  Am.  Jour.  Sci.,  [4],  xxii,  492. 

t  Zeit.  anal.  Chem.,  xxiii,  378. 

t  Am.  Chem.  Jour.,  vii,  373. 

§  F.  A.  Gooch  and  P.  E.  Browning,  Am.  Jour.  Sci.,  [3],  xl,  66. 

II  See  page  457. 


2Q2  METHODS  IN  CHEMICAL  ANALYSIS 

the  iodine  method.  The  conditions  of  the  methods  are  different, 
in  that  in  the  former  the  hydriodic  acid  is  entirely  broken  up 
by  the  action  of  the  arsenic  acid,  and  the  iodine  volatilizes 
easily;  while  in  the  latter  some  hydriodic  acid  must  remain  in 
solution  until  a  very  low  degree  of  concentration  is  reached,  and 
remaining  must  exhibit  its  characteristic  proneness  to  retain  free 
iodine. 

It  is  found  in  practice  that  when  a  solution  made  up  to  contain 
sulphuric  acid,  an  arsenate,  and  potassium  iodide  to  an  amount 
somewhat  in  excess  of  that  theoretically  demanded  to  effect  the 
conversion  of  the  arsenic  acid  to  arsenious  acid,  is  boiled,  iodine 
is  evolved.  The  color  of  the  liquid  passes  from  the  dark  red 
when  the  iodine  is  abundant  through  the  various  gradations  of 
tint  to  a  canary  yellow,  and  then,  as  the  sulphuric  acid  reaches 
a  degree  of  concentration  sufficient  to  determine  by  its  own 
specific  action  the  liberation  of  iodine,  the  color  again  darkens. 
If  the  process  of  concentration  is  continued,  and  much  arsenic 
is  present,  crystals  of  arsenious  iodide  separate  and  form  more 
abundantly  on  cooling.  If  evaporation  is  pushed  still  farther 
arsenious  iodide  begins  to  volatilize,  and  at  the  point  where  the 
sulphuric  acid  fumes  the  liquid  loses  all  color  and  the  arsenic 
has  vanished  more  or  less  completely.  In  one  experiment  con- 
ducted in  this  manner  it  was  found,  by  the  method  to  be  described 
later,  that  of  0.3861  grm.  of  arsenic  pentoxide  originally  present 
with  I  grm.  of  potassium  iodide  and  10  cm.3  of  sulphuric  acid 
[i  :  i]  the  equivalent  of  0.1524  grm.  remained.  In  another  simi- 
lar experiment,  in  which,  however,  only  a  few  milligrams  of  ar- 
senic oxide  were  involved,  not  a  trace  of  arsenic  remained  at  the 
end. 

It  is  obvious  that  two  points  in  this  course  of  action  demand 
attention :  First,  means  must  be  used  for  removing  the  remnant 
of  free  iodine  which  is  withheld  by  the  hydriodic  acid,  or  of 
rendering  it  harmless  in  the  titration  process  to  follow;  and, 
secondly,  the  degree  to  which  the  solution  may  be  concentrated 
without  loss  of  arsenic  must  be  fixed.  In  the  converse  of  this 
process,  the  marked  influence  of  the  amount  of  sulphuric  acid 
present  upon  the  degree  of  concentration  necessary  to  expel  the 
iodine  was  particularly  noted.  In  the  present  case,  the  effect 
of  the  proportion  of  sulphuric  acid  in  solutions  containing  definite 
amounts  of  potassium  iodide  and  potassium  arsenate  is  like- 


ARSENIC,   ANTIMONY  AND  TIN 


293 


wise  of  first  importance.  In  studying  the  effects  of  concentration, 
the  solution  was  made  up  to  about  100  cm.3  and  concentrated  by 
boiling  until  the  color  was  faintest;  then,  to  determine  provi- 
sionally, and  for  preliminary  purposes,  the  point  at  which  vola- 
tilization of  arsenic  was  likely  to  occur,  the  concentration  was 
continued  until  the  arsenious  iodide  began  to  separate.  The 
results  are  tabulated  as  follows: 


KI. 

As205. 

H2S04  [i  :  i]. 

Volume  when 
color  was  lightest. 

Volume  when 
AsI3  appeared. 

grm. 

grm. 

cm.3 

cm.3 

cm.j 

I 

o.  1900 

20 

80 

33 

I 

o.  1900 

15 

65 

25 

I 

o.  1900 

IO 

40 

19 

I 

0.1900 

5 

30 

II 

The  amount  of  sulphuric  acid  which,  considering  rapidity  in 
concentrating  to  the  proper  point,  ease  in  neutralizing  the  acid 
previous  to  titration,  and  general  convenience  in  manipulation, 
seemed  to  be  best  is  10  cm.3  of  the  [i:  i]  mixture.  The  most 
suitable  limit  of  concentration  of  the  solution  appears  to  be 
40  cm.3. 

It  is  manifest  from  the  phenomena  described  that  when  much 
hydriodic  acid  remains  in  the  solution  the  last  portions  of  free 
iodine  cannot  be  completely  removed  by  heat  without  volatiliza- 
tion of  the  arsenic.  It  was  found  that  upon  adding  approximately 
n/ioo  sulphurous  acid  drop  by  drop  to  the  hot  concentrated 
solution  the  point  at  which  the  color  vanished  could  be  deter- 
mined without  difficulty,  but  that  if  the  solution  was  permitted 
to  stand  a  single  minute  the  color  of  iodine  returned,  developed 
by  the  action  of  air  upon  the  hot  hydriodic  acid.  By  diluting 
the  solution  with  cold  water  as  soon  as  the  sulphurous  acid  has 
done  its  work  and  immediately  neutralizing  with  potassium  car- 
bonate, reversion  of  arsenious  acid  to  arsenic  acid  is  precluded, 
magnesia  mixture  producing  in  the  solution  no  precipitate  of  the 
ammonium  magnesium  arsenate. 

The  process  as  recommended  by  Gooch  and  Browning  may  be 
summarized  briefly  as  follows :  To  the  arsenate  taken  in  solu- 
tion in  a  2OO-cm.3  Erlenmeyer  flask  are  added  potassium  iodide 
in  excess  of  the  amount  needed  according  to  the  equation  to 
complete  the  reduction,  and  10  cm.3  of  [1:1]  sulphuric  acid 


294 


^METHODS  IN  CHEMICAL  ANALYSIS 


The  liquid  is  diluted  to  about  100  cm.3  and  boiled  rapidly 
(with  the  precaution  of  trapping  with  a  two-bulbed  tube  hung 
with  the  large  end  downward)*  until  the  volume  diminishes 
to  40  cm.3,  shown  by  a  mark  upon  the  flask.  The  color  of 
free  iodine  is  bleached  by  cautious  additions  of  sulphurous  acid 
(corresponding  roughly  to  centinormal  iodine)  and  the  solution 
is  instantly  diluted  with  water,  nearly  neutralized  with  potas- 
sium carbonate,  and  completely  with  the  acid  carbonate.  The 
liquid  is  cooled  and  titrated  as  usual  with  iodine,  using  starch 
as  an  indicator.  The  whole  operation  is  easily  completed  in  a 
half-hour. 

Results  obtained  by  the  procedure  are  given  in  the  table. 

Reduction  by  Hydriodic  Acid:    Oxidation  by  Iodine  in  Alkaline  Solution. 


KI  taken, 
gnu. 

H2S04  [i  :  i] 
taken. 

cm.j 

As2O6  taken, 
grin. 

As2O5  found, 
grin. 

Error, 
grm. 

i-5 

IO 

0.3861 

0.3862 

+O.OOOI 

i-5 

IO 

0.3862 

0.3856 

—  O.OOo6 

i-S 

IO 

0.3861 

0.3862 

+0.0001 

i-S 

10 

0.3860 

0.3862 

+O.OO02 

.     i-S 

IO 

0.3863 

0.3862 

—  o.oooi 

i-.S 

IO 

0.3862 

0.3862 

o.oooo 

I 

IO 

0.1927 

0.1922 

-0.0005 

I 

IO 

0.1928 

0.1922 

—0.0006 

I 

IO 

0.1930 

0.1925 

—0.0005 

I 

IO 

0.1930 

0.1927 

—0.0003 

I 

IO 

0.1936 

o.  1929 

—  0.0007 

I 

IO 

o.  1929 

0.1928 

—o.oooi 

I 

10 

0.0383 

o  .  0380 

—0.0003 

I 

10 

0.0383 

0.0385 

+O.OOO2 

o-5 

IO 

o  .  0383 

0.0384 

+O.OOOI 

0.4 

IO 

o  .  0383 

0.0385 

+O.O002 

o-3 

IO 

o  .  0383 

0.0386 

+0.0003 

O.2 

IO 

o  .  0383 

0.0384 

+  O.OOOI 

O.2 

10 

0.0076 

o  .  0074 

—  O.O002 

O.2 

IO 

0.0076 

0.0074 

—  0.0002 

0.  2 

IO 

o  .  0038 

0.0034 

—  O.0004 

O.2 

10 

o  .  0038 

0.0034 

—  O.OO04 

1 

In  subsequent  work  by  Gooch  and  Morris  f  it  is  shown  that 
the  process  may  be  shortened  by  restricting  the  volume  at  which 
heating  begins  so  that  the  boiling  need  not  be  extended  beyond 

five  or  six  minutes. 

f 

*  See  Fig.  6,  page  6. 

t  F.  A.  Gooch  and  Julia  C.  Morris,  Am.  Jour.  Sci.,  [4],  x,  151.  . 


ARSENIC,  ANTIMONY  AND  TIN 


295 


According  to  this  slight  modification,  the  solution  of  the  arse- 
nate  is  heated  in  a  {rapped  Erlenmeyer  flask*  with  potassium 
iodide  to  an  amount  about  0.5  grm.  in  excess  of  the  amount 
theoretically  required  and  10  cm.3  of  sulphuric  acid  of  half 
strengthen  a  total  volume  of  50  cm.3  to  75cm.3.  The  liquid 
is  boiled  till  the  iodine  vapors  are  no  longer  visible  in  the  flask 
above  the  liquid,  the  iodine  color  in  the  still  hot  liquid  is 
bleached  by  the  cautious  addition  of  sulphurous  acid,  the  whole 
is  diluted  with  cold  water,  and  cooled  quickly.  The  solution  is 
nearly  neutralized  with  potassium  hydroxide  and  the  neutraliza- 
tion is  completed  with  acid  potassium  carbonate.  The  reduced 
acid  is  titrated  with  iodine  after  adding  the  starch  indicator. 
By  this  procedure  the  results  of  the  following  table  were  obtained. 

Reduction  by  Hydriodic  Acid  and  Sulphurous  Acid:    Oxidation  by  Iodine  in 
Alkaline  Solution. 


Volume. 
cm.1 

H3O3AsO  taken, 
grm. 

HjO3AsO  found, 
grm. 

Error, 
grm. 

35 

0.1559 

0.1559 

o.oooo 

35 

0-1559 

0.1560 

+O.OOOI 

40 

0.1559 

0.1559 

0.0000 

65 

0.1559 

0.1559 

0.0000 

5° 

0.2495 

0.2499 

+0.0004 

50 

0-2557 

0.2449 

—0.0008 

60 

0.3119 

0-3H7 

—  O.OOO2 

60 

0.3119 

0.3120 

+O.OOOI 

75 

0.3119 

0.3124 

+0.0005 

75 

0.3119 

0.3132 

+0.0013 

75 

0.3119 

0.3121 

+O.OOO2 

75 

0.3119 

0-3H5 

—0.0004 

75 

0.3119 

0.3124 

+o  .  0005 

Reduction  by          j/he  process  just  described  for  the  reduction  and 

Hydriodic  Acid:  J 

Titrationof  estimation  of  arsenic  acid,  depending  upon  the  re- 
iodine  Liberated.  movai  by  volatilization  of  all  but  the  last  traces  of 
liberated  iodine,  and  the  conversion  of  this  minute  residue  by 
sulphurous  acid,  involves  no  secondary  reactions  of  a  sort  likely 
to  influence  the  main  effect.  It  is  exact  and  rapid. 

The  method  of  Williamson, f  brought  forward  more  recently, 
depends  upon  the  conversion  of  the  liberated  iodine  to  hydri- 
odic  acid.  The  interaction  at  ordinary  temperatures  of  a  suit- 

*  See  Fig.  6,  page  6. 

f  Jour.  Soc.  Dyers  and  Colorists,  1896,  86-89. 


296  METHODS  IN  CHEMICAL  ANALYSIS 

ably  strong  acid,  hydrochloric  or  sulphuric  acid,  upon  the  mix- 
ture of  the  arsenate  and  iodide,  sets  free  iodine,  and  the  liberated 
iodine  is  converted  to  hydriodic  acid  by  the  action  of  sodium 
thiosulphate,  the  end-point  being  the  disappearance  of  the  iodine 
color. 

According  to  Williamson's  directions,  25-cm.3  portions  of  the 
solution  of  the  arsenate  are  treated  with  potassium  iodide  and 
mixed  with  an  equal  volume  of  hydrochloric  acid  of  sp.  gr. 
1. 1 6.  The  precaution  is  recommended  that  the  strength  of  the 
solution  of  the  arsenate  shall  not  exceed  the  decinormal  value, 
in  order  that  the  dilution  consequent  upon  titration  by  the  thio- 
sulphate may  not  be  too  great;  the  reducing  action  brought 
about  by  the  action  of  the  strong  acid  upon  the  arsenate  and 
iodide  being  reversible  upon  the  dilution  of  liquid  with  water. 
This  procedure  thus  limits  the  process  to  the  determination  of 
about  o.i 8  grm.  of  arsenic  acid  in  25  cm.3  of  the  solution  to  be 
treated  with  an  equal  volume  of  hydrochloric  acid  of  sp.  gr.  1.16. 
Obviously,  however,  the  process  should,  so  far  as  the  reduction 
is  concerned,  be  applicable  to  larger  amounts  of  arsenic,  provided 
the  strength  of  the  acid  is  kept  up  proportionately.  It  is  essen- 
tial that  the  liquid  at  the  end  of  the  titration  should  contain 
approximately  10  per  cent  of  its  mass  of  absolute  hydrochloric 
acid  or  about  one-third  of  its  volume  of  the  aqueous  acid  of 
sp.  gr.  1.16. 

The  arsenic  acid  is  measured  either  by  the  amount  of  stand- 
ard thiosulphate  required  to  bleach  the  iodine  or  by  the  amount 
of  iodine  required  afterward  to  reoxidize  the  arsenious  acid,  after 
neutralizing  with  acid  potassium  carbonate.  If  the  former  alter- 
native is  followed,  the  end  reaction  must  be  the  disappearance 
of  the  yellow  color  of  the  iodine,  since  in  solutions  so  strongly  acid 
it  is  impossible  to  place  dependence  upon  the  starch  indicator; 
in  using  the  latter  alternative,  the  starch  indicator  is,  of  course, 
permissible  and  preferable. 

In  the  direct  titration  of  the  iodine  by  thiosulphate  two  sources 
of  error  present  themselves  as  possibilities:  first,  the  excessive 
liberation  of  iodine  by  the  action  of  air  upon,  the  strongly  acidu- 
lated iodide;  and  second,  the  liability  of  the  thiosulphate,*  if 
present'  even  in  momentary  or  local  excess  during  the  process  of 
titration,  to  break  down  under  the  action  of  strong  acid,  thus 
*  Norton,  see  page  364. 


ARSENIC,  ANTIMONY  AND   TIN 


297 


changing  its  capacity  to  convert  iodine  to  hydriodic  acid.  The 
latter  contingency  should  be  remote  in  proportion  to  the  caution 
used  in  adding  the  thiosulphate  and  in  keeping  the  liquid  well 
stirred;  the  former  must  of  necessity  vary  with  the  acidity  of 
the  solution  containing  the  iodide,  the  time  of  exposure  to  atmos- 
pheric action,  and  the  degree  of  contact  with  the  air  incidental 
to  stirring.  How  far  each  of  these  possibilities  is  likely  to  inter- 
fere in  the  practical  conduct  of  an  ordinary  analysis  has  been  in- 
vestigated by  Gooch  and  Morris.* 

The  effects  likely  to  result  simply  from  the  strong  acidifica- 
tion of  the  solution  containing  potassium  iodide,  and  their  varia- 
tion for  conditions  of  dilution  representing  the  beginning  and 
the  end  of  a  titration  on  the  lines  laid  down,  are  shown  in  the 
following  table.  The  solution  of  potassium  iodide  was  diluted 
as  indicated  before  the  addition  of  the  acid,  and  the  iodine  set 
free  was  titrated  by  thiosulphate. 

Effect  of  Concentration  of  Acid  and  Time  of  Action  upon  Potassium  Iodide. 


HC1 
(sp.  gr. 
1.16)  taken. 

cm.8 

KI  taken, 
giro. 

Total 
volume. 

cm.3 

Na2S203  added 
at  once.     In  terms 
of  H303AsO. 

grm. 

Na2S2O3  added 
after  5  minutes. 
In  terms  of 
H303AsO. 

grin. 

Na2S2O3  added 
after  stirring 
5  minutes.    In 
terms  of 
H3O3AsO. 

grm. 

2C 

2 

CO 

0.0013 

2< 

2 

7C 

O  .  0004 

2C 

2 

50 

0.0035 

25 

2 

75 

0.0019 

2C 

2 

co 

O  0042 

25 

2 

vc 

O   OO2I 

50 
50 

CO 

2 
2 

2 

IOO 

150 

IOO 

0.0017 
o  .  0004 

o  .  003  ^ 

50 

2 

ISO 

0.0019 

co 

2 

IOO 

O   OO3  C. 

CO 

2 

ISO 

o  0014 

The  concentration  of  acid  and  the  time  before  titration  are, 
obviously,  the  essential  factors.  The  absolute  amount  of  acid 
present  and  the  stirring  seem  to  make  little  difference. 

As  to  the  action  of  the  hydrochloric  acid  on  small  amounts 
of  the  thiosulphate,  there  is  the  evidence  of  the  experiments  de- 
tailed in  the  following  statements,  in  which  I  cm.3,  2  cm.3  and  5 

*  F.  A.  Gooch  and  Julia  C.  Morris,  Am.  Jour.  Sci.,  [4],  x,  151. 


298 


METHODS  IN  CHEMICAL  ANALYSIS 


Effect  of  Concentration  of  Acid  upon  Thio sulphate. 


Iodine  to 

Error  of 

Iodine  to 

Error  of 

HC1 

(sp.  gr. 

Volume 
before 

Na2S2O3 
nearly  K/IO.      In 

color  with- 
out dilution. 

.  titration 
without 

color  after 
diluting 

titration 
after 

1.16). 

titration. 

terms  of  H3O3AsO. 

In  terms  of 
H303AsO. 

dilution. 
In  terms  of 
H303AsO. 

to  75  cm.3 
In  terms  of 
H3O3AsO. 

dilution. 
In  terms  of 
H303AsO. 

cm.8 

cm.8 

cm.3           grm. 

grm. 

grin. 

grm. 

grm. 

25 

26 

I 

0.0071 

0.0062 

—0.0009 

0.0071 

O  .  OOOO 

25 

50 

I 

0.0071 

0.0071 

O  .  OOOO 

0.0071 

0  .  OOOO 

25* 

50 

I 

0.0071 

0.0079 

+0.0008 

o  .  0079 

+0.0008 

50 

2 

0.0141 

0.0146 

+0.0005 

0.0146 

+0.0005 

25* 

50 

2 

0.0141 

0.0157 

+0.0016 

0.0157 

+0.0016 

25 

30 

5 

0.0353 

0.0336 

—  0.0017 

0.0374 

+0.0024 

25 

50 

5 

0.0353 

0-0359 

+0.0006 

0-0359 

+o  .  0006 

25* 

50 

5 

0.0353 

0.0411 

+0.0058 

0.0411 

+0.0058 

In  these  experiments  the  acid  stood  in  contact  with  the  thiosulphate  5  minutes  before  titration. 


cm.3  of  nearly  n/io  thiosulphate  were  exposed  to  the  action  of  25 
cm.3  hydrochloric  acid  (sp.  gr.  1. 1 6),  without  dilution  or  diluted 
with  an  equal  volume  of  water,  and  titrated  with  nearly  n/io 
iodine.  The  condition  of  acidity  when  the  volume  of  50  cm.3 
contains  25  cm.3  of  hydrochloric  acid  (sp.  gr.  1.16)  is  that  of  the 
beginning  of  titration  of  Williamson's  process.  In  order  that 
the  effect  of  error  due  to  such  action  upon  the  determination  of 
arsenic  acid  may  appear  immediately,  the  thiosulphate  and  iodine 
used  are  expressed  in  terms  of  that  acid. 

• 
Time  Effect  of  Acid  upon  Thiosulphate. 


HC1 
(sp.  gr. 
1.16). 

cm.3 

KI. 

grm. 

Volume. 
cm.8 

Na2S203 
nearly  «/io. 
In  terms  of 
H3O3AsO. 

cm.1           grm. 

Iodine  in 
terms  of 
H3O.,AsO, 
at  once. 

grm. 

Iodine  in 
terms  of 
H303AsO, 
after  5  min. 

grm. 

Error  in 
terms  of 
H303AsO. 

grm. 

25 
25 
25 
25 
25 
25 

25 
25 
25 
25 
25 
25 

2 
2 
2 
2 
2 
2 

2 
2 
2 
2 
2 
2 

50 
75 
50 
75 
So 
75 

50 
75- 
50 
75 
50 
75 

I 
I 
2 
2 
5 

5 

I 

I 

2 

2 

5 
5 

0.0071 
0.0071 
0.0141 
0.0141 
0.0353 
0-0353 

0.0071 
0.0071 
0.0141 
0.0141 
0-0353 
0.0353 

0.0057 
0.0071 
0.0131 
0.0143 
O  .  03  2  2 
0.0357 

—  0.0014 
O.OOOO 
—  O.OOIO 
+0.0002 
—  O.OO2I 
+O.OOO4 

-0.0043 
—  O.OO04 
—  O.O025 
—  O.OOO2 
—O.OO4I 
+0.0008 

.  . 

O.OO28 
0.0067 
0.0116 
0.0139 
0.0314 
0.0361 

ARSENIC,  ANTIMONY  AND   TIN 


299 


The  two  sources  of  error  due  to  the  action  of  hydrochloric 
acid,  the  liberation  of  iodine  and  the  decomposition  of  the  thio- 
sulphate,  naturally  tend  to  neutralize  one  another,  but  the 
completeness  of  such  neutralization  must  be  largely  a  matter  of 
chance  in  the  varying  conditions  of  actual  analysis.  The  experi- 
ments of  the  preceding  table,  in  which  n/io  thiosulphate,  to  the 
amount  of  I  cm.3,  2  cm.3  and  5  cm.3,  was  added  to  the  liquid, 
50  cm.3  and  75  cm.3,  containing  25  cm.3  acid,  and  titrated  with 
iodine  at  once,  and  after  five  minutes,  were  made  to  test  the 
effects  for  the  conditions  of  dilution  prevailing  at  the  beginning 
and  at  the  end  of  a  titration. 

It  is  clear  that  under  the  conditions  covered  by  the  experi- 
ments of  the  two  preceding  tables  the  decomposition  of  the 
thiosulphate  is  likely  to  occur  in  greater  or  less  degree,  and  that 
when  the  acid  of  sp.  gr.  1.16  is  not  much  diluted  the  products 
of  decomposition  are  not  oxidized  by  the  iodine  completely. 
The  latter  observation  is  quite  in  harmony  with  the  fact  that 
sulphur  dioxide  bleaches  iodine  in  strong  hydrochloric  acid  only 
slowly  and  incompletely.  In  such  cases  dilution  favors  further 
action  of  the  iodine,  but  results  obtained  by  titration  with  iodine 
in  the  acid  solution  diluted  with  an  equal  amount  of  water  are 
unmodified  by  further  dilution. 

In  the  following  tables  are  recorded  actual  determinations  of 
arsenic  according  to  Williamson's  process.  To  each  25  cm.3  of 
the  arsenate  were  added  I,  2  or  3  grm.  of  potassium  iodide  and 
25  cm.3  hydrochloric  acid  (sp.  gr.  1.16).  The  iodine  was  bleached 

Williamson's  Procedure. 


HC1. 

KI. 

Volume  at 
beginning 
of  titration. 

Volume  at 
end  of 
titration. 

H2KAsO4 
in  terms  of 
H3O3AsO. 

H303AsO 

found. 

Error. 

cm.8 

grm. 

cm.3 

cm.3 

grm. 

grm. 

grm. 

25 

2 

50 

51 

O.CX362 

o  .  0085 

+0.0023 

25 

2 

50 

52 

0.0125 

0.0156 

+0.0031 

25 

2 

SO 

55 

0.0312 

0.0350 

+0.0038 

25 

2 

50 

55 

0.0624 

0  .  0666 

+0.0042 

2-5 

2 

50 

73 

0.1559 

0.1588 

+0.0029 

25 

2 

50 

73 

0-1559 

0.1587 

+0.0028 

25 

2 

50 

73 

0.1559 

O.I5QI 

+0.0032 

25 

2 

50 

73 

O.I5S9 

0.1595 

+0.0036 

25 

3 

50 

73 

0.1559 

0-1595 

+0.0036 

25 

i 

50 

73 

0.1559 

0.1581 

+O.OO22 

25 

2 

50 

73 

0.1559 

0.1581 

+0.0022 

25 

2 

50 

73 

0.1559 

0.1588 

+0.0029 

300 


METHODS  IN  CHEMICAL  ANALYSIS 


by  nearly  decinormal  thiosulphate  without  addition  of  the  starch 
indicator,  which  loses  all  delicacy  in  the  presence  of  strong  acid. 
The  time  occupied  by  each  titration  was  about  five  minutes. 
The  standards  of  the  arsenate  were  determined  by  the  vapori- 
zation process,*  the  purity  of  reagents  employed  in  that  process 
having  been  proved  by  trying  the  process  in  the  estimation  of  a 
solution  of  arsenic  acid  made  by  oxidizing  pure  decinormal  arse- 
nious  acid  by  iodine. 

The  range  of  error  in  these  results  is  from  +0.0023  grm.  to 
+0.0042  grm.  with  a  mean  of  +0.0031  grm.  —  not  very  different 
from  what  might  be  expected  from  the  effect  of  the  interaction 
of  the  strong  hydrochloric  acid  and  the  iodide  alone.  The 
counter  effect  due  to  the  decomposition  of  the  thiosulphate  is 
not  large,  yet  it  is  probably  real,  as  will  appear  in  the  sequel. 

In  the  following  series  of  determinations,  made  with  new 
solutions  and  new  standards  throughout,  the  arsenic  acid  was 
determined  in  two  ways:  (I)  The  iodine  set  free  by  25  cm.3 
of  hydrochloric  acid  (sp.  gr.  1.16)  and  3  grm.  potassium  iodide, 
the  solution  having  a  total  volume  of  50  cm.3  at  beginning  and 
of  75  cm.3  at  the  end,  was  titrated  by  sodium  thiosulphate; 
(II)  The  arsenious  acid  remaining  after  the  first  titration  by 
sodium  thiosulphate  was  titrated,  after  being  neutralized  with 
a.cid  potassium  carbonate  by  iodine,  in  the  presence  of  the  starch 
indicator. 

Titration  of  Iodine  Liberated  and  Titration  of  Ar senile  Produced. 


HjKAs04  taken 
in  terms  of 
HsOsAsO. 

HAAsO  found 
by  the  thiosulphate. 

Error. 

H3O3AsO  found 
by  titration  of 
H3O3As  with  iodine. 

Error. 

grm. 

grm 

grm. 

grm. 

grm. 

0.1767 

0.1708 

+0.0031 

0.1776 

+0-0009 

0.1767 

o.  1708 

+0.0031 

0.1777 

+O.OOi:o 

0.1767 

0.179$ 

+O.OO28 

0.1785 

+O.OOl8 

0.1767 

0.179? 

+0.0026 

0.1785 

+0.0018 

0.1767 

0.1701- 

+0.0027 

0.1780 

+0.0013 

0.1767 

0.1798 

+0.0031 

0.1785 

+o.ooc8 

The  average  error  of  the  first  operation  is  0.0029  grm.,  not 
far  from  that  of  the  previous  series;  the  error  of  the  second 
operation,  the  titration  of  the  arsenious  acid,  amounts  on  the 
average  to  0.0014  grm-  I*1  tne  second  operation  the  error  due 

*  See  page  291. 


ARSENIC,  ANTIMONY  AND  TIN  301 

to  over  use  of  the  thiosulphate  by  iodine  set  free  outside  the  main 
reaction  is  obviously  eliminated.  The  tetrathionate  present 
after  neutralization  with  acid  potassium  carbonate  is  unaffected 
by  iodine,  as  was  found  by  titrating  25  cm.3  of  n/io  iodine 
mixed  with  25  cm.3  hydrochloric  acid  (sp.  gr.  1.16)  by  the  thio- 
sulphate, neutralizing  with  acid  potassium  carbonate,*  adding 
starch  and  getting  the  starch  blue  with  a  single  drop  of  n/io 
iodine.  The  average  error  of  this  process  (0.0014),  therefore,  is 
probably  due  to  the  products  of  decomposition  of  the  thiosul- 
phate in  the  first  operation. 

From  the  foregoing  experiments  it  is  clear  that  an  arbitrary- 
correction  of  about  0.0030  grm.  must  be  deducted  from  the 
indications  of  Williamson's  process  of  direct  titration  by  thiosul- 
phate, made  with  the  greatest  care  under  the  conditions  men- 
tioned ;  and  that  a  correction  varying  from  one-half  that  amount 
(0.0015  §rm-)  to  nothing  (according  to  the  amount  of  arsenious 
acid  present),  when  the  determination  is  made  by  iodine  after 
neutralization  with  acid  potassium  carbonate.  After  making 
these  arbitrary  corrections  in  the  results  of  the  preceding  table,, 
the  individual  variations  fall  within  reasonable  limits. 

On  the  other  hand,  the  vaporization  process,  in  which  the 
arsenate  is  reduced  by  boiling  with  sulphuric  acid  and  potas- 
sium iodide  in  the  manner  described,!  gives  indications  reasonably 
regular  and  accurate  without  the  application  of  an  arbitrary 
correction. 

The  Detection  and  Approximative  Estimation  of  Minute  Quantities 
of  Arsenic  in  Copper. 

Sanger's  successful  application  of  the  Berzelius-Marsh  process 
to  the  quantitative  determination  of  arsenic  in  wall  papers  and 
fabrics,|  by  the  comparison  of  test  mirrors  with  standard  mirrors 
carefully  prepared  under  the  conditions  of  the  test,  opens  the 
way,  naturally,  to  the  similar  estimation  of  minute  amounts  of 

*  It  is  worthy  of  note  that,  as  was  found,  it  is  not  possible  to  substitute  an 
alkali  hydroxide  for  the  carbonate  in  the  early  stages  of  the  process  of  neutral- 
ization, on  account  of  the  decomposing  effect  of  the  former  reagent  upon  the 
tetrathionate.  This  effect  is  in  proportion  to  the  heating  of  the  solution,  but 
is  never  wholly  absent  even  when  ice  is  intermixed  with  the  liquid  and  the 
greatest  care  taken  to  prevent  a  rise  of  temperature. 

t  See  pages  29 1,  294. 

J  Am.  Chem.  Jour.,  xiii,  431. 


302  METHODS  IN  CHEMICAL  ANALYSIS 

arsenic  in  any  substances  which  may  be  submitted  to  the  process 
immediately  or  after  suitable  preparation.  Gooch  and  Moseley  * 
have  studied  the  application  of  Sanger's  process  to  the  deter- 
mination of  traces  of  arsenic  in  copper. 

It  has  been  shown  by  Headden  and  Sadler  f  that  the  presence 
of  copper  in  the  Marsh  generator  is  instrumental  in  holding  back 
the  arsenic.  It  is  obvious,  therefore,  that  means  must  be  em- 
ployed for  the  complete  removal  of  the  copper  from  the  arsenic 
before  the  solution  of  the  latter  is  put  into  the  reduction  flask, 
and  there  is  no  method  by  which  arsenic  may  be  removed  from 
copper  easily,  and  without  loss,  aside  from  those  methods  which 
depend  upon  the  volatility  of  arsenious  chloride  from  solution  in 
strong  hydrochloric  acid.  Of  such  methods,  on  the  score  of 
rapidity  in  execution,  accessibility  of  pure  materials,  and  com- 
pactness of  apparatus,  preference  is  to  be  given  to  that  process 
which  is  based  upon  the  simultaneous  action  of  strong  hydro- 
chloric acid  and  potassium  bromide  upon  the  salt  of  arsenic.  J 

To  get  the  copper  into  condition  for  the  application  of  the 
process  of  separation  from  arsenic,  it  is  sufficient  to  dissolve  an 
amount  not  exceeding  I  grm.  in  nitric  acid  somewhat  diluted  with 
water,  to  add  to  the  solution  2  cm3  or  3  cm.3  of  strong  sulphuric 
acid,  and  to  evaporate  the  liquid  until  fumes  of  the  sulphuric 
acid  are  disengaged  abundantly.  A  single  treatment  of  this  sort 
serves  to  remove  the  nitric  acid  so  completely  that  no  interfer- 
ence with  the  normal  action  of  the  Marsh  apparatus  is  apparent 
in  the  subsequent  operation.  The  residue  after  concentration 
is  diluted  with  water  to  about  5  cm.3  and  washed  into  the  dis- 
tillation flask  with  an  amount  of  the  strongest  hydrochloric  acid 
(sp.  gr.  1. 20)  equal  to  that  of  the  remainder  of  the  liquid.  It  is 
desirable  that  the  entire  volume  of  the  liquid  should  not  much 
exceed  10  cm.3.  The  flask,  which  has  a  capacity  of  40  or  50  cm.3, 
is  inclined  at  an  angle  of  about  45°  and  joined  by  means  of  a 
pure  rubber  stopper  to  a  bent  pipette  which  serves  as  a  distilla- 
tion tube.  The  lower  end  of  the  vertical  limb  of  the  pipette 
dips  beneath  the  surface  of  about  5  cm.3  of  hydrochloric  acid  of 
half-strength  contained  in  a  test  tube  which  is  cooled  and  sup- 
ported by  water  nearly  filling  an  Erlenmeyer  flask.  A  gram  of 

*  F.  A.  Gooch  and  H.  P.  Moseley,  Am.  Jour.  Sci.,  [3],  xlviii,  292. 
t  Am.  Chem.  Jour.,  vii,  342. 
t  See  page  316. 


ARSENIC,  ANTIMONY  AND   TIN  303 

potassium  bromide  is  introduced,  and  the  distillation  (which  may 
be  completed  in  three  or  four  minutes)  is  pushed  nearly  to  dry- 
ness.  The  flask  is  washed  out,  another  portion  of  potassium 
bromide  is  introduced,  and  the  first  distillate  is  introduced  and 
redistilled  as  before,  excepting  that  the  condensation  is  this  time 
effected  in  pure  water.  This  second  operation  serves  merely  to 
hold  back  traces  of  copper  carried  over  in  the  first  distillation, 
but  the  addition  of  the  potassium  bromide  in  the  second  distil- 
lation is  quite  as  necessary  as  in  the  first,  since  the  bromine 
liberated  in  the  process  has  the  effect  of  reoxidizing  the  arsenic 
in  the  receiver  and  so  making  that  element  nonvolatile  under  the 
conditions  until  the  reducing  agent  is  again  introduced.  The 
free  bromine  in  the  final  distillate  must  be  reconverted  to  hydro- 
bromic  acid  before  the  contents  of  the  receiver  may  be  introduced 
into  the  reduction  flask,  and  this  effect  may  be  most  easily  and 
unobjectionably  accomplished  by  the  addition  of  a  little  stannous 
chloride  dissolved  in  hydrochloric  acid  of  half-strength  and  puri- 
fied from  arsenic  by  prolonged  boiling.  Incidentally  and  simul- 
taneously the  arsenic  is  reduced  to  the  arsenious  form,  and, 
though  Sanger  has  shown  that  minute  amounts  of  arsenic  are 
completely  eliminated  from  the  solution  in  the  reduction  flask 
when  that  element  is  introduced  in  the  higher  form  of  oxidation, 
it  is  our  experience  that  the  rapidity  of  elimination  of  the  arsenic 
is  so  increased  by  the  introduction  of  the  small  amount  of  stan- 
nous chloride  needed  to  bleach  the  bromine  that  the  mirror 
appears  in  from  five  to  ten  minutes  and  is  practically  complete 
in  half  an  hour,  especially  if  the  precaution  is  taken  to  add  a  little 
more  stannous  chloride,  according  to  Schmidt's  suggestion,*  after 
the  operation  has  been  in  progress  about  twenty  minutes. 

Schmidt  has  shown  that  the  addition  of  stannous  chloride  to 
the  Marsh  apparatus  in  action  not  only  does  not  effect  the  reten- 
tion of  arsenic,  as  many  other  metallic  salts  do,  but  actually 
brings  about  the  final  evolution  in  the  form  of  the  hydride  of 
that  portion  of  the  arsenic  which  may  have  been  deposited  during 
the  process  in  elementary  form  upon  the  zinc. 

The  Sanger  apparatus  is  used  in  form  essentially  unchanged; 

but  the  zinc  in  the  reserve  generator  is  coated  with  copper  by  the 

action  of  a  solution  of  copper  sulphate,  since  in  this  way  it  is 

made  more  sensitive  to  the  action  of  the  dilute  sulphuric  acid, 

*  Zeit.  anorg.  Chem.,  i,  353. 


304 


METHODS  IN  CHEMICAL  ANALYSIS 


while  the  presence  of  copper  (which  is  of  course  out  of  the  ques- 
tion in  the  reduction  flask)  can  be  of  no  disadvantage  in  the 
reserve  generator  and  may  even  serve  a  useful  end  in  fixing 
traces  of  arsenic  if  the  zinc  and  acid  employed  are  not  absolutely 
free  from  that  element.  In  the  formation  of  the  mirror,  too,  it 
has  proved  to  be  an  advantage  to  inclose  the  portion  of  the  glass 
tube  to  be  heated  in  a  short  thin  tube  of  iron  or  nickel  slightly 
larger  than  the  glass  tube  and  kept  from  contact  with  it  except 
at  the  ends,  which  are  notched  and  bent  inward.  By  keeping 
the  outer  tube  of  metal  at  a  low  red  heat  it  is  possible  to  diminish 
the  tendency  of  the  arsenic*  particularly  when  the  amounts  are 
fairly  large,  to  form  a  double  mirror  corresponding  to  the  allo- 
tropic  conditions  of  the  arsenic.  It  is  necessary,  moreover,  to 
substitute  hydrochloric  acid  for  the  sulphuric  acid  usually  em- 
ployed in  the  reduction  flask;  but,  though  the  opinion  is  current 
that  hydrochloric  acid  introduces  difficulties  in  the  Marsh  test, 
no  evidence  of  the  formation  of  a  zinc  mirror  in  the  ignition  tube 
or  of  other  unfavorable  action  due  to  the  use  of  pure  hydrochloric 
acid  has  been  noted  in  this  operation.  It  is,  of  course,  obvi- 
ous that  the  hydrochloric  acid  used  must  be  arsenic  free. 

Tests  of  this  process  were  made  with  copper-prepared  by  elec- 
trolyzing  in  ammoniacal  solution  the  purest  copper  sulphate 
obtainable  and  stopping  the  deposition  before  the  solution  had 
become  exhausted.  This  copper,  in  which  no  arsenic  was  found, 
was  dissolved  in  nitric  acid,  arsenic  in  the  higher  condition  of 
oxidation  was  added,  and  the  process  of  the  separation  of  the 
arsenic  from  the  copper  and  conversion  to  the  mirror  carried  out 
in  the  manner  described. 

A  rsenic  in  Copper. 


Copper  taken. 

Arsenic  taken. 

Mirror  estimated  (by 
comparison  with 
standard  mirror). 

Error. 

gnu. 

mgrm. 

mgrm. 

mgrm. 

None. 

None. 

None. 

None. 

0.7 

None. 

None. 

None. 

o-5 
0.5 

0.005 

O.OII 

0.003 
0.013 

—  0.002 

+  O.002 

0-35 
0.3 

O.O2O 

0.030 

0.015 
0.030 

—0.005 
None. 

0.43 

0.040 

0.035 

—0.005 

0.44 

0.050 

0.040 

—  o.oio 

ARSENIC,   ANTIMONY  AND   TIN 


305 


It  is  plain  from  the  results  given  above  that  the  method  is 
capable  of  detecting  sharply  minute  amounts  of  arsenic  in  copper 
and  of  effecting  the  estimation  of  quantities  less  than  0.05  mgrm. 
with  some  approximation  to  accuracy. 

There  is,  as  Sanger  has  pointed  out,  a  good  deal  of  variation 
even  in  standard  mirrors  made  with  all  possible  care  and  pre- 
caution, and  in  the  estimation  of  mirrors  containing  as  much 
as  0.05  mgrm.  of  arsenic  the  uncertainty  of  comparison  as  well  as 
the  actual  variation  of  the  mirror  is  considerable. 

When  a  sample  of  copper  is  under  test  which  may  contain 
more  than  0.05  mgrm.  of  arsenic,  it  is  desirable  to  introduce  into 
the  reduction  flask  the  measured  solution  containing  the  arsenic 
gradually  and  in  definite  portions,  and  to  judge  by  the  formation 
of  the  mirror  in  an  interval  of  ten  minutes  after  the  introduction 
of  a  portion  of  this  test  solution  whether  it  is  wiser  to  add  the 
entire  solution  or  to  estimate  the  arsenic  in  the  entire  solution 
from  that  found  in  an  aliquot  portion. 

Results  of  the  analysis  of  several  samples  of  commercial 
electrolytic  copper  are  appended.  The  last  two  represented, 
presumably,  the  very  purest  electrolytically  refined  copper 
obtainable  commercially. 

Arsenic  in  Commercial  Copper. 


Copper  taken, 
giro. 

Arsenic  found, 
mgrrn. 

Percentage  of 
arsenic. 

Sample  A  

o.  3 

O.OI<J 

O  OOS 

Sample  B  

0.3 

0.030 

O  OIO 

Sample  C 

{' 

O.OlS 

O.OOlS 

Sample  D          .    . 

1  I 
1 

0.015 
0.005 

0.0015 
0.0005 

1  I 

0.005 

0.0005 

The  Separation  of  Arsenic  from  Copper  by  Precipitation  as  Am- 
monium Magnesium  Ar senate. 

Arsenic  acid  in  an  alkali  salt  may  be  compbtely  precipitated 
as  the  ammonium  magnesium  arsenate,  even  in  presence  of  am- 
monium salts,  by  adding  the  solution,  with  stirring,  to  a  sufficient 
excess  of  magnesia  mixture  kept  ammoniacal.* 

*  See  page  288. 


306  METHODS  IN  CHEMICAL  ANALYSIS 

The  fact  that  ammonium  magnesium  arsenate  is  insoluble 
in  presence  of  an  abundance  of  magnesia  mixture,  while  many 
salts  of  copper  are  soluble  in  ammonia,  suggests  that  arsenic 
existing  as  an  arsenate  may  be  separated  from  copper  in  ammo- 
niacal  solution  by  magnesia  mixture.  Obviously,  the  action  of 
the  ammoniacal  copper  solution  on  cellulose  renders  it  impossi- 
ble to  make  such  an  estimation  by  the  use  of  paper  niters ;  but 
ammonium  magnesium  arsenate,  as  has  been  shown,  may  be 
filtered  off  upon  a  mat  of  fine  asbestos  under  pressure,  in  a  per- 
forated platinum  crucible. 

Gooch  and  Phelps*  have  found  that,  while  a  single  precipita- 
tion is  sufficient  to  effect  the  separation  of  small  amounts  of  the 
arsenic  acid  from  copper,  copper  is  held  in  appreciable  amount, 
apparently  in  combination,  when  the  amount  of  the  magnesium 
pyroarsenate  exceeds  a  few  milligrams.  By  dissolving  the  first 
precipitate  and  reprecipitating,  reasonably  good  separations  may, 
however,  be  effected  for  amounts  not  exceeding  0.4  grm.  For 
larger  amounts  a  second  dissolving  and  reprecipitation  must  be 
made. 

According  to  procedure  outlined,  the  slightly  acid  solution  of 
the  alkali  arsenate  is  run  from  a  burette  with  stirring  into  the 
slightly  ammoniacal  solution  of  copper  sulphate  and  magnesia 
mixture,!  and  the  mixture  is  made  distinctly  ammoniacal.  The 
precipitate  is  transferred  to  a  weighed  crucible,  and,  after  rinsing 
once  with  distilled  water  made  faintly  ammoniacal,  is  dissolved 
in  hot  hydrochloric  acid  [i  :  3].  For  convenience  in  handling  the 
solutions  the  filtrate  is  received  in  a  beaker  under  an  evacuated 
bell-jar  rather  than  in  the  usual  filter  flask.  After  cooling  and 
adding  ammonia  nearly  to  neutrality,  the  solution  is  poured 
with  stirring  into  an  abundance  of  magnesia  mixture  kept  con- 
stantly ammoniacal.  The  precipitate  obtained  in  this  way  is 
collected  on  the  asbestos  felt  used  for  the  first  filtration,  the  first 
portion  of  the  filtrate  being  employed  in  each  case  to  remove 
the  last  portions  of  the  ammonium  magnesium  arsenate  from 
the  platinum  dish.  For  amounts  exceeding  0.4  grm.  the  process 
of  dissolving,  reprecipitating  and  filtering  is  repeated.  After 
rinsing  all  traces  of  reagents  from  the  precipitate  with  distilled 
water  made  faintly  ammoniacal  (20  cm.3~5O  cm.3),  the  crucible 

*  F.  A.  Gooch  and  M.  A.  Phelps,  Am.  Jour.  Sci.,  [4],  xxii,  488. 
t  See  page  289. 


ARSENIC,  ANTIMONY   AND  TIN 


307 


and  contents  are  dried  over  a  low  Bunsen  flame  until  all  ammonia 
is  driven  off,  and  then  ignited  cautiously.  Results  of  this  pro- 
cedure are  given  in  the  table. 


Separation  of  Arsenic  from  Copper. 


CuSO4. 
grm. 

H2KAs04. 
cm.3 

Magnesia 
Mixture 

cm.8 

Mg2As2O7. 

Error  in 
teims  of 
arsenic. 

grm. 

Theory, 
grm. 

Found, 
grm. 

Error, 
grm. 

One  precipitation. 


2 

O.  2 

25 

0.0015 

O.OO15 

0  .  OOOO 

o.oooo 

2 

O.  2 

25 

0.0015 

0.0015 

O  .  OOOO 

o.oooo 

2 

O.  2 

25 

0.0015 

0.0015 

o.oooo 

o.oooo 

2 

I 

25 

0.0077 

0.0086 

+o  .  0009 

+0.0004 

Two  precipitations. 


2 

0.2 

25-25 

0.0015 

0.0012 

—0.0003 

—  O.OOOI 

2 

0.2       . 

25-25 

0.0015 

o  .  0009 

—0.0006 

—  0.0003 

2 

I 

25-25 

0.0077 

0.0072 

—0.0005 

—  O.OOO2 

2 

I 

25-25 

0.0077 

0.0079 

+O.OOO2 

+O.OOOI 

2 

5 

25-25 

0.0386 

0.0389 

+0.0003 

+0.0001 

2 

5 

25-25 

0.0386 

0.0383 

—  0.0003 

—  O.OOOI 

2 

10 

25-25 

0.0766 

0.0754 

—  O.OOI2 

—0.0006 

2 

10 

25-25 

0.0766 

0.0754 

—  O.OOI2 

—  0.0006 

2 

25 

25-25 

0.1931 

0.1927 

—  O.OOO4 

—  O.OOO2 

2 

25 

1OO-IOO 

0.1931 

o.  1919 

—  O.OO12 

—  O.O006 

2 

50 

50-50 

0.3862 

0.3867 

+0.0005 

+  O.OOO2 

2 

50 

25-25 

0.3862 

0.3864 

+0.0002 

+0.0001 

Three  precipitations. 


2 

50 

50-  50-50 

0.3830 

0.3822 

—0.0008 

—  0.0004 

2 

IOO 

50-  50-50 

0.7660 

0.7653 

—0.0007 

—0.0003 

2 

100 

50-  50-50 

o.  7660 

o.  7656 

—  0.0004 

—  O.OOO2 

2 

IOO 

100-100-100 

0.7724  I    0.7726 

+  O.OOO2 

+O.OOOI 

To  recover  amounts  of  the  arsenate  so  small  as  not  to  be  at 
once  predpitable  by  magnesia  mixture  (at  the  outset  or  after 
the  solution  in  hydrochloric  acid),  the  solution,  best  contained 
in  a  platinum  dish,  is  frozen  in  a  mixture  of  ice  and  salt.  Upon 
allowing  the  frozen  mass  to  melt  at  the  ordinary  room  temper- 
ature, insoluble  magnesium  arsenate  remains.  Results  of  the 
procedure  are  given  in  the  table  below. 


308 


METHODS   IN   CHEMICAL  ANALYSIS 


Recovery  of  Small  Amounts  by  Freezing. 


Mg2As207. 

CuSO4 

"M"          '  t 

NH4C1 

Theory. 

Found. 

Error. 

grin. 

cm.  3 

grm. 

grm. 

grm. 

grm. 

2 

25-25 

0.0015 

0.0015 

o  .  oooo 

2 

25-25 

0.0015 

0.0015 

o.oooo 

2 

25-25 

IO 

0.0015 

O.OOI2 

—0.0003 

The  lodometric  Determination  of  Antimonic  Acid  and  of  Antimonic 
Acid  and  Arsenic  Acid. 

Bunsen's  method  of  determining  qualitatively  the  condition 
of  oxidation  of  salts  of  antimony,  by  boiling  these  substances  in 
solution  with  potassium  iodide  and  hydrochloric  acid  and  noting 
whether  the  liquid  takes  the  color  of  free  iodine,  has  been  applied 
successfully  to  the  quantitative  determination  of  antimony  in 
its  highest  condition  of  oxidation.  Weller*  distils  the  iodine 
from  the  solution,  collects  it  in  the  distillate  and,  determining  it 
volumetrically,  calculates  from  the  amount  found  the  antimonic 
salt  which  has  set  it  free  according  to  the  equation 

SbCl5  +  2  HI  =  SbCl3  +  2  HC1  +  I2. 

The  advantage  of  treating  the  residue,  rather  than  the  distillate, 
in  analytical  processes  in  general  which  involve  distillation  is, 
however,  so  obvious  that  Gooch  and  Gruener  f  have  determined 
the  conditions  under  which  Bunsen's  reaction  may  be  applied  in 
such  manner  that  the  antimony  is  held  and  estimated  directly 
in  the  residue.  The  general  plan  of  work  was  laid  down  in  a 
similar  process  elaborated  in  this  laboratory  for  the  reduction 
of  arsenic  acid-t  According  to  this  process,  the  arsenic  to  be  re- 
duced is  taken  in  a  solution  of  appropriate  dilution,  and  treated 
with  sulphuric  acid  in  adjusted  amount  and  an  excess  of  potassium 
iodide;  the  liquid  thus  prepared  is  boiled  to  a  definite  degree  of 
concentration,  the  iodine  then  remaining  unexpelled,  if  any,  is 
bleached  by  the  very  careful  addition  of  dilute  (centinormal) 

*  Ann.  Chem.,  ccxiii,  246. 

t  F.  A.  Gooch  and  H.  W.  Gruener,  Am.  Jour.  Sci.,  [3],  xlii,  213. 

{  See  page  291. 


ARSENIC,  ANTIMONY  AND   TIN  309 

sulphurous  acid,  and  the  liquid  is  immediately  diluted  and  neu- 
tralized ;  after  cooling,  the  reduced  arsenic  is  titrated  by  standard 
iodine  in  presence  of  starch. 

It  was  found  in  preliminary  experimentation  that  the  same 
general  plan  of  treatment  is  available  in  the  handling  of  anti- 
monic  compounds,  but  it  is  necessary  to  take  precautions  to 
prevent  the  deposition  of  the  antimony  from  solution  upon  the 
addition  of  the  sulphuric  acid.  Tartaric  acid  accomplishes  this 
effect  satisfactorily  and  does  not,  as  the  result  proved,  introduce 
undesirable  complications.  It  appears  that  the  dilution  of  the 
solution  at  which  the  crystalline  iodide  or  oxyiodide  separates  out 
during  the  boiling  is  greater  than  is  the  case  when  similar  amounts 
of  arsenic  are  dealt  with,  and  that  concentration  to  45  cm.3  is 
sufficient  to  cause  crystallization  and  slight  sublimation  when  the 
amount  of  antimonious  oxide  present  (with  excess  of  potassium 
iodide  and  10  cm.3  of  sulphuric  acid,  I  :  i)  is  approximately 
0.2  grm.  Otherwise  the  process  as  employed  in  the  reduction 
of  arsenic  appears  to  be  applicable  to  the  similar  treatment  of 
antimony. 

According  to  the  procedure  developed  experimentally,  the 
antimonate,  in  amount  not  exceeding  the  equivalent  of  O.-2  grm. 
of  antimonious  oxide,  is  treated  in  a  trapped  3OO-cm.3  Erlen- 
meyer  flask*  with  tartaric  acid  (4  grm.)  and  sulphuric  acid  to 
acidity  and  thereafter  with  10  cm.3  of  [i  :  i]  sulphuric  acid  and 
i  grm.  of  potassium  iodide.  The  mixture  is  boiled,  after  intro- 
ducing folded  platinum  foil,  to  prevent  bumoing  of  the  liquid, 
and  a  trap  to  prevent  mechanical  loss.  When  the  liquid  has 
been  concentrated  to  a  volume  of  45  cm.3  to  55  cm.3  the  boiling 
is  stopped,  the  color  bleached  by  the  cautious  addition  of 
sulphurous  acid  (approximately  centinormal).  The  solution  is 
diluted,  nearly  neutralized  with  sodium  hydroxide,  made  alkaline 
by  hydrogen  sodium  carbonate  in  an  excess  amounting  to  about 
20  cm.3  of  the  saturated  solution,  and  titrated  with  the  standard 
(decinormal)  iodine  after  the  addition  of  starch. 

Less  concentration  in  the  boiling  may  result  in  incomplete  re- 
duction and  greater  concentration  when  the  amount  of  antimony 
present  is  large  may  result  in  loss  of  that  element  by  volatilization. 

Results  of  experiments  made  under  the  conditions  prescribed 
are  given  in  the  following  table . 

*  See  Fig.  6,  page  6. 


3io 


METHODS  IN  CHEMICAL  ANALYSIS 


Determination  of  Antimonic  Acid. 


Final 
volume. 

Tartar  emetic, 
taken. 

Sb2O3  taken 
and  oxidized. 

Iodine  used 
in  final 
oxidation. 

Sb2O3  found. 

Error. 

cm.* 

grm. 

grm. 

cm.a 

grm. 

grm. 

55 

0.5023 

0.2178 

0.3827 

0.2178 

0.0000 

55 

0.5015 

0.2175 

0.3806 

O.  2166 

—  O.OOO9 

So 

0.5007 

o.  2172 

0.3814 

0.2171 

—  O.OOOI 

So 

o  .  5039 

0.2185 

0.3839 

0.2185 

O.OOOO 

45 

0.5001 

o.  2169 

0.3818 

0.2173 

+  0.0004 

45 

0.5004 

o.  2170 

0.3825 

o.  2176 

-|-O.OOO6 

The  antimonate  used  in  these  experiments  was  made  by  titrat- 
ing tartar  emetic  in  solution  in  presence  of  acid  sodium  carbonate 
by  n/io  iodine.  This  treatment  of  the  tartar  emetic  served  the 
double  purpose  of  providing  a  perfectly  definite  antimonic  salt 
and  of  adjusting  the  standard  of  the  iodine  to  be  used  subse- 
quently in  reoxidizing  the  antimony  after  its  reduction.  The 
imperfection  of  the  process,  whatever  it  may  be,  whether  in 
the  reduction  or  elsewhere,  becomes  apparent  and  is  measured 
immediately  by  the  difference  between  the  amounts  of  iodine 
employed  in  the  two  oxidations.* 

The  method  is  also  applicable  to  the  reduction  and  estimation 
of  antimony  and  arsenic  in  association,  as  is  shown  by  the  fol- 
lowing test  results. 

Determination  of  Antimonic  Acid  and  Arsenic  Acid. 


Difference 

Error  in  terms  of 

Final 
volume. 

Tartar 
emetic 
taken. 

Sb,0» 
taken 
and 
oxi- 
dized. 

As,08 
taken 
and 
oxi- 
dized. 

Iodine 
used  in 
first  oxi- 
dation.* 

Iodine 
used-  in 
final  oxi- 
dation. 

between 
the  amounts 
of  iodine 
used  in  the 
two  oxi- 

Sb203. 

As203. 

dations. 

cm.* 

gnu* 

grm. 

grm. 

cm.3 

cm.8 

cm.3 

grm. 

grm. 

50 

o  .  1530 

0.0870 

0.0500 

10-37 

19-43 

+0.06 

+0  .  0004 

—0.0003 

50 

0.1503 

0.0855 

o  .  0405 

T9-05 

19.02 

—  0.03 

—  O.OOO2 

—  O.OOOI 

50 

0.1503 

0.0855 

o  .  0544 

20.05 

19.97 

-0.08 

—  O.OOO6 

—0.0004 

50 

0.1503 

0.0855 

0.0495 

IQ.05 

19.00 

—  0.05 

—  0.0004 

+0.0003 

*  To  form  the  antimonate  and  arsenate. 

*  Hale  has  shown  that  a  variation  between  the  value  of  the  iodine  deter- 
mined against  tartar  emetic  and  the  value  determined  against  arsenious  oxide 
is  due  to  variation  in  the  composition  of  the  tartar  emetic  and  is  not  inherent 
in  the  process  of  titration  (see  page  40). 


ARSENIC,  ANTIMONY  AND   TIN  311 

The  Separation  of  Antimony  from  Arsenic  by  the  Simultaneous 

Action  of  Hydrochloric  Acid  and  Hydriodic  Acid,  and 

the  Estimation  of  Antimony  in  the  Residue. 

The  separation  of  arsenic  from  antimony  by  taking  advan- 
tage of,  the  difference  in  volatility  of  the  lower  chlorides  is  due 
to  Fischer.*  This  method  of  treatment  consists  in  the  reduc- 
tion of  the  chlorides  by  means  of  ferrous  chloride  and  the  vola- 
tilization of  the  arsenic  by  repeated  distillations  of  the  mixture 
with  hydrochloric  acid  of  20  per  cent  strength  added  in  succes- 
sive portions.  The  process  has  been  subsequently  modified  by 
Huf schmidt t  by  the  substitution  of  gaseous  hydrochloric  acid, 
introduced  in  a  continuous  current  into  the  distilling  mixture,  for 
the  aqueous  acid,  and  later  changed  further  and  improved  by 
Classen  and  Ludwig,t  who  employ  ferrous  sulphate,  or  ammonio- 
ferrous  sulphate,  in  place  of  the  less  easily  prepared  ferrous 
chloride.  In  its  latest  form  the  method  is  exceedingly  exact, 
but  the  conditions  are  such  that  the  antimony  in  the  residue 
must  be  determined  gravimetrically.  Gooch  and  Danner  §  have 
arranged  the  process  so  that  the  determination  of  the  antimony 
may  be  made  by  a  rapid  volumetric  method,  by  substituting 
for  the  iron  salt,  which  utterly  precludes  the  direct  volumetric 
estimation  of  the  antimony,  another  reducer  —  hydriodic  acid  — 
which  can  interfere  in  no  way  with  the  subsequent  determination 
of  the  antimony  by  the  well-known  iodometric  method. 

It  has  been  shown  that  arsenic ||  and  antimony**  may  be  re- 
duced by  the  action  of  hydriodic  acid  applied  under  conditions 
such  that  the  arsenic  shall  not  volatilize.  In  the  present  case 
the  reducing  action  of  hydriodic  acid  takes  place  in  the  presence 
of  strong  hydrochloric  acid,  and  at  the  boiling  temperature  of  the 
solution,  —  conditions  arranged  to  bring  about  the  volatilization 
of  the  arsenic  as  rapidly  as  possible. 

The  method  of  proceeding  is  briefly  summarized  in  the  follow- 
ing statement :  To  the  solution  of  the  oxides  of  arsenic  and  anti- 
mony, taken  in  amounts  not  exceeding  0.5  grm.  of  each,  potas- 

*  Ann.  Chem.,  ccviii,  182. 
f  Ber.  Dtsch.  chem.  Ges.,  xvii,  2245. 
J  Ber.  Dtsch.  chem.  Ges.,  xviii,  mo. 

§  F.  A.  Gooch  and  E.  W.  Danner,  Am.  Jour.  Sci.,  [3],  xlii,  308. 
II  See  page  291. 
**  See  page  308. 


3I2 


METHODS   IN  CHEMICAL  ANALYSIS 


sium  iodide  is  added  in  a  little  more  than  the  equivalent  quantity, 
and  enough  strong  hydrochloric  acid  to  raise  the  entire  volume  of 
the  solution  to  100  cm.3.  The  solution  is  saturated  with  hydro- 
chloric acid  gas  and  submitted  to  distillation  in  a  current  of  that 
gas  until  the  volume  decreases  to  50  cm.3  or  a  little  less.  The 
liquid  is  cooled  rapidly,  treated  first  with  an  excess  of  sulphurous 
acid  and  then  with  iodine  to  the  exact  oxidation  of  the  former 
reagent;  and,  after  the  addition  of  I  grm.  of  tartaric  acid  to 
every  0.2  grm.  of  antimonious  oxide,  the  acid  present  is  nearly 
neutralized  with  sodium  hydroxide.  The  solution  is  made 
alkaline  with  hydrogen  sodium  carbonate  added  in  excess  to 
an  amount  corresponding  to  10  cm.3  of  the  saturated  solution 
for  every  o.i  grm.  of  antimonious  oxide  present.  Titration  with 
decinormal  iodine  gives  the  antimony  quickly  and  with  a  fair 
degree  of  accuracy.  The  whole  process  requires  about  an  hour 
and  a  half  for  completion.  Results  obtained  according  to  the 
procedure  are  given  below. 

Separation  of  Antimony  from  Arsenic. 


H2K- 
As04 
taken. 

KI 

taken. 

Volume. 

Color. 

Sbo03 
taken. 

Sb203 
found. 

Error. 

Initial. 

Final. 

On  cooling. 

With 

grm. 

grm. 

cm.3 

cm.* 

grm. 

grm. 

grm. 

o-S 

o-S 

100 

50 

Pale  yellow. 

Faint. 

0.2268 

O.  2265 

—  0.0003 

0-5 

o-5 

TOO 

50 

Pale  yellow. 

Faint. 

0.2306 

0.2300 

—  0.0006 

0-5 

o-S 

IOO 

50 

Pale  yellow. 

Faint. 

0.2272 

O.  2264 

—  0.0008 

When  the  treatment  with  sulphurous  acid  is  omitted  the  final 
titration  indicates  apparent  deficiency  in  the  antimony  oxide 
amounting  sometimes  to  more  than  two  milligrams.  This  defi- 
ciency appears  to  be  due,  at  least  in  part,  to  the  presence,  at 
the  time  of  neutralization,  of  a  small  amount  of  iodine  chloride, 
which  might  easily  be  produced  by  the  oxidizing  effect  of  anti- 
monic  and  arsenic  oxides  upon  the  mixed  halogen  acids. 

The  Detection  of  Arsenic,  and  of  Antimony  with  Tin,  in  Mixtures 
Containing  Compounds  of  These  Elements. 

Upon  the  well-known  fact  that  hot  strong  hydrochloric  acid 
dissolves  easily  the  sulphides  of  antimony  and  tin,  but  arsenious 
sulphide  to  only  a  slight  degree,  is  based  the  simplest  and  most 
rapid  method  in  common  use  for  the  separation  of  arsenic  from 


ARSENIC,  ANTIMONY  AND   TIN  313 

antimony  and  tin.  Unfortunately,  however,  the  forcible  treat- 
ment necessary  to  bring  about  the  solution  of  large  amounts  of 
antimony  is  sufficient  *  to  dissolve  small  quantities  of  arsenious 
sulphide,  so  that  for  the  purposes  of  general  analysis  the  method 
is  inadequate.  Koehler  f  has  shown  that  only  the  arsenic  is 
precipitated,  and  that  very  completely,  when  hydrogen  sulphide 
acts  upon  the  solution  of  arsenious  and  antimonious  salts  in 
hydrochloric  acid  of  20  per  cent  strength,  but  the  adaptability  of 
Koehler's  treatment  to  the  detection  of  arsenic  in  the  ordinary 
course  of  analysis  is  limited  by  the  necessity  of  so  constituting 
the  solution  to  be  tested  that  hydrogen  sulphide  shall  occasion  no 
deposit  of  free  sulphur  to  conceal  or  be  mistaken  for  a  precipita- 
tion of  arsenious  sulphide.  In  the  course  of  analysis  the  mixed 
sulphides  of  arsenic,  antimony,  and  tin,  remaining  after  the 
removal  of  the  sulphides  insoluble  in  alkaline  sulphides  and 
recovered  from  solution  by  the  action  of  hydrochloric  acid, 
require  for  their  complete  solution  the  action  of  an  oxidizing 
agent,  which  must,  of  course,  interfere  with  the  immediate  use 
of  Koehler's  method. 

Action  of  Gooch  and  Hodge  J  have  applied  the  action  of  hy- 

Ac^and'potes-  drochloric  aci£l  anci  potassium  iodide  to  the  reduc- 
siumiodide.  tion  of  the  salts  of  arsenic,  antimony,  and  tin,  with 
the  volatilization  of  the.  arsenic  by  repeated  distillations.  The 
arsenic  may  be  looked  for  in  the  distillate  by  hydrogen  sulphide 
by  Koehler's  method  after  removal  of  free  iodine  by  means  of 
stannous  chloride.  To  effect  the  repeated  distillations  of  small 
portions  of  concentrated  hydrochloric  acid  upon  mixtures  of  the 
salts  with  potassium  iodide,  an  apparatus  which  is  essentially 
the  distillation  apparatus  of  Mohr  is  employed.  This  consists 
of  a  25-cm.3  flask  fitted  by  means  of  a  rubber  stopper  to  a  pipette, 
bent,  drawn  out  at  the  lower  end,  and  dipped  into  a  test  tube 
which  is  at  the  same  time  supported  and  cooled  in  a  flask  partly 
filled  with  water.  The  pipette  tube  is  wide  enough  (about 
0.7  cm.  in  diameter)  to  prevent  the  formation  of  bubbles  within 
it,  and  the  bulb,  holding  about  20  cm.3,  is  sufficiently  large  to 
retain  any  liquid  which  may  be  momentarily  forced  back  by  the 
accidental  cooling  of  the  flask  during  the  distillation. 

*  Rose-Finkener,  Anal.  Chem.,  ii,  423. 

t  Zeit.  anal.  Chem.,  xxix,  192. 

t  F.  A.  Gooch  and  B.  Hodge,  Am.  Jour.  Sci.,  [3],  xlvii,  382, 


METHODS  IN  CHEMICAL  ANALYSIS 


The  mixture  of  salts  of  arsenic,  antimony,  and  tin,  in  the  higher 
condition  of  oxidation,  is  put  in  the  flask  with  3  grm.  of  potas- 
sium iodide  in  5  cm.3  of  water  and  5  cm.3  of  the  strongest  hydro- 
chloric acid  (sp.  gr.  1.20).  The  distillation  is  carried  nearly  to 
dryness,  and  the  distillate  is  condensed  in  10  cm.3  of  a  mixture 
of  strong  hydrochloric  acid  and  water  in  equal  parts.  The 
iodine  evolved  during  the  distillation  is  bleached  by  the  addition 
to  the  distillate  of  stannous  chloride  dissolved  in  hydrochloric 
acid  of  half -strength,  and  hydrogen  sulphide  is  passed  to  pre- 
cipitate the  arsenic  if  present.  The  residue  in  the  flask  is  treated 
with  10  cm.3  of  the  strongest  hydrochloric 
acid  and  the  process  of  distillation  repeated, 
but  this  time  the  distillate  is  condensed  in 
10  cm.3  of  water  in  order  that  the  final 
acidity  of  the  liquid  may  be  that  of  acid 
of  half-strength,  and  so,  after  bleaching  by 
stannous  chloride,  immediately  available 
for  the  test  -for  arsenic  by  hydrogen  sul- 
phide. Subsequent  treatments  of  the  resi- 
due are  carried  out  similarly  until  arsenic 
ceases  to  appear  in  the  distillate. 

Four  distillations  of  io-cm.3  portions  of 
the  strongest  acid  suffice  to  transfer  o.oi 
grm.  of  arsenic  completely  to  the  distillate,  while  a  single  dis- 
tillation appears  to  be  sufficient  to  volatilize  anything  less  than 
0.003  grm. 

Experiments  made  with  antimony  taken  in  the  form  of  purified 
tartar  emetic  and  oxidized  by  iodine  in  alkaline  solution  previous 
to  treatment,  either  alone  or  with  arsenic,  show  that  antimony  is 
discoverable  in  the  residues  when  even  so  little  as  o.oooi  grm.  of 
that  element  is  originally  introduced,  though  it  is  evident  that  a 
portion  of  the  antimony  may  pass  to  the  distillate  when  much 
of  it  is  present  in  the  flask.  Indeed,  when  large  quantities  of 
antimony  are  treated  the  appearance  of  the  brownish-red  fumes 
of  antimonious  iodide  in  the  distilling  tube  may  serve  as  an 
indication  that  the  concentration  should  go  no  further,  since  the 
antimonious  iodide  will,  if  it  reaches  the  receiver  in  quantity, 
impart  to  the  distillate  a  color  which  is  not  discharged  by  the 
stannous  chloride  used  to  bleach  the  iodine.  In  this  case  it 


Fig.  23. 


ARSENIC,   ANTIMONY  AND   TIN 


315 


will  be  necessary  to  look  subsequently  for  a  precipitate  of 
arsenious  sulphide  in  a  liquid  of  its  own  tint.  The  amount  of" 
antimony  volatilized  seems  to  be  proportioned  to  the  amount 
present,  and,  when  the  distillation  is  properly  conducted,  enough 
antimony  remains  in  the  residue  to  be  found  if  it  was  originally 
present  in  discoverable  quantity. 

The  results  of  similar  work  with  tin  alone,  and  with  tin  and 
arsenic,  go  to  show  that,  though  tin  like  antimony  may  pass  to 
the  distillate  under  the  conditions,  enough  tin  always  remains  to 
be  found  in  the  residue,  if  the  amount  was  originally  appreciable. 

Tests  with  Hydrochloric  Acid  and  Potassium  Iodide. 


Arsenic  taken 
as  HsOsAsO. 

grm. 

Antimony 
taken  as 
H303SbO. 

grm. 

Tin  taken  as 
SnCl4. 

grm. 

Precipitation  by 
H2S  in  successive 
distillates. 

Precipitation  by 
H2S  in  the  residue 
dissolved  in  water. 

O   OOOI 

(  I.    Found.           ) 

None 

o  0033 

I  II.    None.           ) 
j  I.    Found.           ) 

None. 

0.0050 

(  II.    None.           J 
\  I-III.    Found.   | 

None. 

O.OIOO 

1  IV.    None.          J 
j  I-IV.    Found.    ) 

None. 

O.  IOOO 

7  V.    None.           J 
j  I-VII.  Found.  ) 

None. 

O    OOOI 

{  VIII.    None.      J 
I.   None 

Distinct  color 

o  .  0050 

O  OOOI 

i  I-IV.   Found.    ) 

Distinct  color 

O.OOOI 

0.4 

7  V.    None.            \ 
(  I.    Found.           \ 

Large 

O.OIOO 

0.4 

O  OOOI 

7  II.    None.           J 
j  I-IV.    Found.    ) 
fy.   None.           { 
I     None 

Large. 
Distinct  color 

O.OIOO 

O   OOOI 

(  I-IV.   Found.    \ 

Distinct  color 

O.OOOI 

0.0005 

I  V.    None.           J 
J  I.    Found.           ) 

Distinct 

O   OIOO 

o  0005 

(  II.    None.           j 
(  I-IV.    Found,    j 

Distinct 

O.OOOI 

o  ? 

I  V.    None.           J 
(  I.    Found.           ) 

Lar^e 

O.OIOO 

o.  s 

1  II.    None.           J 
(  I-IV.   Found.    / 

Large 

7  V.   None.           J 

A  single  distillation,  which  may  easily  be  completed  in  five 
minutes,  is  sufficient  to  discover  the  presence  of  o.oooi  grm.  of 
arsenic  associated  with  so  much  as  0.4  grm.  or  0.5  grm.  of  anti- 


31 6  METHODS  IN  CHEMICAL  ANALYSIS 

mony  or  tin,  and  to  remove  from  the  residue  amounts  of  arsenic 
not  exceeding  0.003  grm.  When  larger  amounts  of  arsenic  are 
to  be  removed,  so  that  the  tin  and  antimony  may  be  obtained 
free  from  that  element,  the  result  may  be  accomplished  by 
a  suitable  number  of  distillations;  or,  inasmuch  as  only  a  little 
iodine  remains  after  the  first  distillation,  the  end  may  be  at- 
tained by  dissolving  the  residue  in  hydrochloric  acid  of  half- 
strength,  bleaching  the  iodine  with  exactly  the  necessary  amount 
of  sulphurous  acid  or  sodium  thiosulphate  (since  the  use  of  the 
stannous  chloride  is  here  precluded),  and  passing  hydrogen 
sulphide. 

The  details  of  experimental  tests  are  given  in  the  preceding 
statement. 

Action  of  Gooch  and  Phelps*  have  shown  that  the  action  of 

Hydrochloric       hydrobromic  acid  upon  arsenic  acid  is  so  similar  to 

Acid  and  .  .  .  t 

Potassium  that  or  hydnoaic  acid  that  potassium  bromide  may 
Bromide.  with  advantage  be  substituted  for  the  iodide  in  the 

process  of  reduction  and  distillation.  According  to  the  procedure 
indicated,  the  mixture  to  be  analyzed  is  introduced  into  the  flask f 
with  5  cm.3  of  water,  3  grm.  of  potassium  bromide  and  5  cm.3  of 
hydrochloric  acid  of  full  strength  (sp.  gr.  1.20).  The  end  of  the 
pipette  tube  is  dipped  into  5  cm.3  of  hydrochloric  acid  of  half- 
strength  contained  in  this  test  tube  used  as  a  receiver,  and  the 
distillation  is  carried  on  until  the  liquid  in  the  flask  has  almost 
entirely  passed  to  the  receiver.  The  residue  is  treated  with 
10  cm.3  of  the  strongest  hydrochloric  acid,  and  the  distillation  is 
repeated  with  the  modification  that  this  time  the  condensation 
is  effected  by  passing  the  volatile  material  into  10  cm.3  of  water, 
so  that  the  liquid  in  the  receiver  at  the  end  of  this  operation  may 
have  the  acidity  of  hydrochloric  acid  of  half -strength.  This 
process  of  treating  the  residue  with  the  strongest  hydrochloric 
acid  and  distilling  is  continued  until  arsenic  ceases  to  be  dis- 
coverable in  this  distillate.  At  the  beginning  of  the  distillation 
bromine  is  liberated  and  collects  in  this  distillate;  but  later,  as 
the  arsenious  chloride  volatilizes  and  condenses  again,  the  color 
of  the  bromine  in  the  distillate  vanishes  with  the  simultaneous 
reconversion  of  the  arsenic  to  the  higher  form  of  oxidation.  In 
such  a  solution,  especially  if  it  is  not  very  hot,  hydrogen  sulphide 

*  F.  A.  Gooch  and  I.  K.  Phelps,  Am.  Jour.  Sci.,  [3!,  xlviii,  216. 
f  See  Fig.  23,  page  314.  . 


ARSENIC,   ANTIMONY  AND   TIN 


31? 


precipitates  the  arsenic  only  slowly,  but  the  addition  of  a  little 
stannous  chloride  dissolved  in  hydrochloric  acid  of  half-strength 
to  the  hot  solution  reduces  the  arsenic  to  the  lower  form  of  oxida- 
tion and  prepares  the  way  for  the  immediate  precipitation  of 
arsenious  sulphide  by  hydrogen  sulphide.  Antimonic  acid  is 
likewise  reduced  under  the  conditions  of  the  distillation;  but,  as 
Koehler  has  shown,*  neither  the  small  amounts  of  antimony  and 
tin,  which  if  present  originally  may  pass  partially  to  the  distillate> 
nor  the  tin  added  later  to  effect  the  reduction  of  the  arsenic, 
will  be  precipitated  by  hydrogen  sulphide  under  the  existing 
conditions  of  temperature  and  acidity. 

The  results  of  experiments  are  recorded  in  the  accompanying 
table. 

Tests  with  Hydrochloric  Acid  and  Potassium  Bromide. 


Arsenic  taken 
as  H3As04. 

grm. 

Antimony 
•    taken  as 
H3SbO«. 

grm. 

Tin  taken  as 
SnCl4. 

grm. 

Precipitation  by  H2S 
in  successive  distillates 
after  treatment  with 
SnCl2. 

Precipitation  by  HjS 
in  the  residue  dis- 
solved in  water. 

I     None 

None 

I-X     None 

Fuint  coloration 

O   OOOI 

f  I.   Found.          ) 

None 

O.OOIO 

1  II.   None.          f 
(  I.    Found.          ( 

None 

O   OIOO 

(  II.   None.          j 
j  I-II.   Found.     / 

None 

O    IOOO 

|  III.   None.        ( 
j  MIL   Found,  i 

Fciint  coloration 

0.4000 

{  IV.    None.         f 
j  I-VI.   Found     ( 

Orange 

I  .0000 

)VII.    None.       J 
j  I-X.   Found.     / 

precipitation.* 
Orange 

O  4000 

1  XI.    None.         J 
I     None 

precipitation.* 
Larpe 

O.OOOI 

0.4000 

i  I-II.    Found.     ) 

Larc'e 

O  .OOOI 

O.OOOI 

I  III.    None.        J 
j  I.   Found.          ) 

Distinct  color 

O   OOIO 

O   OOOI 

1  II.    None.          { 
i  I.   Found.          i 

Distinct  color 

O.OIOO 

O.OOOI 

1  II.    None.          ) 
(  I-II.   Found.     ) 

Distinct  orange 

o  4000 

1  III.   None.        J 
I     None 

O   OOOI 

o  4000 

i  I-II.   Found.     I 

T  .ftrat* 

O.OOOI 

O.OOOI 

I  III.   None.        J 
t  I.   Found.          J 

Distinct  color 

1  II.   None.          ( 

*  Subsequently  identified  as  antimony  sulphide  by  depositing  the  metal  on  platinum. 
*  Zeit.  anal.  Chem.,  xxix,  192. 


318  METHODS  IN  CHEMICAL  ANALYSIS 

The  lodometric  Determination  of  Arsenic  and  Antimony,  and  of 
Associated  Copper. 

The  determination  of  arsenic  and  antimony  in  the  filtrate 
from  cuprous  iodide  after  the  titration  of  free  iodine  by  sodium 
thiosulphate  has  been  studied  by  Heath*  who  has  thus  made  use 
of  processes  described  elsewhere  f  for  the  reduction  of  arsenic 
acid  and  antimonic  acid  by  the  action  of  potassium  iodide  and 
sulphuric  acid.  The  reactions  involved  in  these  processes  may 
be  expressed  by  the  general  symbols, 

M2O5  +  4  HI  =  M2O3  +  H2O  +  2  I2 
and  M2O3  +  2  I2+  2  K2O  =  M2O5  +  4  KI. 

Trials  of  the  method  were  made  with  tartar  emetic  and  potas- 
sium arsenate.  The  tartar  emetic  was  dissolved  in  water  and 
the  antimony  oxidized  to  the  higher  condition  by  means  of 
standard  iodine  solution  in  presence  of  sodium  or  potassium 
bicarbonate,  the  amount  of  iodine  required  being  taken  as  a 
measure  of  the  amount  of  antimony  used.  The  solution  of 
antimony  thus  obtained,  or  the  solution  of  arsenic  taken  as 
potassium  arsenate,  was  acidified  and  a  known  volume  of  a 
standard  solution  of  copper  nitrate  was  added.  The  copper  was 
determined  iodometrically  with  precautions  recommended  else- 
where, t 

During  the  determination  of  copper,  mineral  acids  must  not 
be  present  on  account  of  their  tendency  to  bring  about  reduction 
of  the  higher  salts  of  arsenic  and  antimony  by  action  of  the  excess 
of  potassium  iodide  used  in  throwing  out  the  copper.  Such 
action  causes  high  results  on  copper  and  low  results  on  arsenic 
and  antimony.  A  mixture  of  acetic  acid  and  potassium  iodide 
reduces  the  higher  salts  slowly.  Tartaric  acid  is  somewhat 
irregular  in  action  and  tends  to  cause  an  interfering  precipi- 
tation of  acid  potassium  tartrate.  The  action  of  citric  acid  is, 
however,  satisfactory  if  the  solution  is  not  allowed  to  stand.  The 
experimental  results  show  that  in  the  iodometric  determination  of 
copper  associated  with  arsenic  there  must  be  no  delay  in  titrat- 
ing the  iodine.  Antimonic  acid  is  not  reduced  appreciably,  in  a 
reasonable  time,  under  similar  conditions. 

*  F.  H.  Heath,  Am.  Jour.  Sci.,  [4],  xxv,  513. 
t  See  pages  29 1,  308. 
I  See  pages  118,  123. 


ARSENIC,  ANTIMONY  AND  TIN  319 

In  the  course  of  preliminary  work  the  fact  developed  that 
tetrathionic  acid,  which  results  from  titration  of  free  iodine  by 
sodium  thiosulphate  in  the  copper  determination,  makes  trouble 
in  the  subsequent  operation.  For,  when  the  solution  is  boiled, 
after  addition  of  sulphuric  acid,  the  tetrathionic  acid  decomposes 
to  give  hydrogen  sulphide  and  free  sulphur,  and  sulphides  of 
antimony  and  arsenic  may  be  precipitated.  It  was  found,  how- 
ever, that  if  liquid  bromine  is  added  to  the  cold  solution  in  suffi- 
cient quantity  to  decompose  all  the  excess  of  potassium  iodide 
present,  and  the  solution  then  boiled,  there  is  very  little  subse- 
quent trouble  on  account  of  tetrathionic  acid. 

The  procedure  for  the  determination  of  copper  and  arsenic 
or  of  copper  and  antimony  may  be  outlined  as  follows :  — To 
the  solution  containing  the  copper  and  also  the  arsenic  or  anti- 
mony in  the  higher  condition  of  oxidation,  add  I  grm.  to  2  grm. 
of  citric  acid.  To  precipitate  amounts  of  copper  not  exceeding 
0.3  grm.  in  a  volume  of  50  cm.3,  add  3  grm.  of  potassium  iodide; 
in  a  volume  of  100  cm.3,  add  5  grm.  of  potassium  iodide.  Titrate 
the  free  iodine  with  n/io  sodium  thiosulphate.  The  reactions 
involved  in  the  copper  determination  are 

2  CuSO4  +  4  KI  ->  2  K2SO4  +  Cu2I2  +  I2 
and 

I2  +  2  Na2S2O3  =  2  Nal  +  Na2S4O6. 

Filter  off  the  cuprous  iodide  on  asbestos.  To  the  filtrate  add 
I  cm.3  of  liquid  bromine  and  boil  the  solution  in  an  Erlenmeyer 
flask,  using  a  trap  to  prevent  loss  by  spattering.  If,  after 
boiling  for  a  short  time  and  allowing  the  large  amount  of  free 
iodine  to  volatilize,  the  solution  does  not  become  clear,  cool  it, 
add  a  little  more  bromine  (0.5  cm.3)  and  boil  again.  When  the 
solution  has  become  clear,  concentrate  it  somewhat  (to  about 
60  cm.3)  to  expel  excess  of  bromine.  Dilute  to  about  100  cm.3, 
add  2  grm.  of  potassium  iodide  and  boil  to  a  volume  of  50  cm.3 
Cool  the  solution,  bleach  the  free  iodine  by  adding  sulphurous 
acid,*  using  starch  as  indicator.  Dilute  to  100  cm.3,  add  n/io 
iodine  to  color  and  just  bleach  by  careful  addition  of  dilute  sul- 
phurous acid  from  a  pipette.  Neutralize  the  solution  with  sodium 
or  potassium  bicarbonate  and  titrate  the  arsenic  or  antimony 
with  standard  iodine  solution  in  the  usual  way. 

*  Compare  page  295. 


320  METHODS  IN  CHEMICAL  ANALYSIS 

Following  are  tables  showing  results  obtained  by  this  procedure 
Copper  and  Arsenic. 


KI 

used. 

giro. 

Volume 
at  end  of 
precipi- 
tation. 

cm.3 

Cu 

taken. 

grm. 

Cu 
found. 

grm. 

Error 
inCu. 

grm. 

Liquid 
bromine 
used. 

cm.3 

KI 

used. 

grm. 

As 
taken. 

grm. 

As 
found. 

grm. 

Error 
in  As. 

grm. 

3 

SO 

0.0700 

o  .  0700 

O.OOOO 

I  .0 

2 

0.1238 

0.1231 

—  O.OOO/ 

3 

50 

0.0700 

0.0693 

—  O.OOO7 

I.O 

2 

0.1238 

0.1231 

—  0.0007 

3 

50 

0.0875 

0.0869 

—  O.OOO6 

1  .0 

2 

0.1238 

0.1239 

+0.0001 

4 

50 

o  .  0700 

0.0698 

—  O.OOO2 

(0.61 

(0.4  [ 

I 

0.1238 

0.1235 

-0.0003 

4 

45 

0.0700 

0.0703 

+O.O003 

i  .0 

2 

0.1238 

0.1247 

+o  .  0009 

4 

50 

0.0700 

O.C700 

0.0000 

i  .0 

2 

0.1238 

0.1235 

-0.0003 

5 

70 

o.  1400 

0.1407 

+0.0007 

(I.O) 

(0.5  ) 

2 

0.1238 

0.1233 

+o  .  0005 

5 

SO 

0.0910 

0.0907 

—0.0003 

I.O 

(  0.5  J 

2 

0.1238 

0.1239 

+O.OOOI 

4 

So 

0.0700 

0.0703 

+0.0003 

1  .0 

2 

0.1238 

0.1237 

—  o.oooi 

4 

50 

0.0875 

0.0879 

+0.0004 

1  .0 

2 

0.1238 

0.1234 

—0.0004 

5 

65 

o.  1400 

O.I4IO 

+O.OOIO 

I.O 

3 

0.1238 

0.1239 

+O.OOOI 

3 

30 

0.0875 

0.0860 

—0.0015 

0.8 

2 

0.0495 

0.0493 

—  O.OOO2 

Copper  and  Antimony. 


KI 
used. 

grm. 

Volume 
atendol 
precipi- 
tation. 

cm.3 

Cu 
taken. 

grm. 

Cu 
found. 

grm. 

Error 
in  Cu. 

grm. 

Liquid 
bromine 
used. 

cm.3 

KI 
used. 

grm. 

Sb 
taken. 

grm. 

Sb 
found. 

grm. 

Error 
inSb. 

grm. 

4 

85 

0.0700 

0.0703 

+0.0003 

(I.O) 

(0.5  ) 

2 

0.1417 

o.  1421 

+0.0004 

4 

80 

0.0700 

0.0701 

+O.OOOI 

I.O 

I  0.5  J 

2 

0.1727 

0.1725 

—  O.OOO2 

4 

80 

0.0875 

0.0869 

—0.0006 

(I.O 

1  0.5  ( 

2 

0.1286 

o.  1289 

+0.0003 

5 

80 

0.0735 

0.0739 

+0  .  0004 

I.O 

1  0.5  i 

2 

0.1641 

0.1645 

+o  .  0004 

5 

90 

0.1050 

0.1050 

0.0000 

I.O 

1  0.5  f 

2 

0.1378 

0.1372 

—0.0006 

4 

75 

0.0875 

0.0874 

—  O.OOOI 

(I.O) 

{0.5  ( 

2 

0.1329 

0.1326 

—0.0003 

6 

"5 

o  .  0700 

0.0707 

+o  .  0007 

1.3 

(0.5  f 

2 

0.2474 

0.2477 

+0.0003 

8 

1  20 

0-1575 

0.1571 

—0.0004 

Ull 

2 

0.1419 

0.1413 

—0.0006 

From  the  results  obtained  it  seems  possible  by  this  method 
to  separate  and  determine  copper  and  arsenic,  or  copper  and 
antimony,  with  errors  of  only  a  few  tenths  of  a  milligram.  It 
is  also  possible  to  determine  the  sum  of  arsenic  and  antimony 


ARSENIC,  ANTIMONY  AND   TIN 


321 


5          0           CO          ^0          ° 

00 

CO 

NO 

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M 

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CN 

0       o       o       o       o 

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8  gji;     g 

o       o       o       o       o 

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M*^*4-              M 

o       o       o      o      o 

o 

o 

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1      1      1      1      1 

1 

1 

1 

I 

1 

£&  -• 

O            ^"          *>•          ON           O 

8 

CO 

a 

NO 

8- 

w^'d     g 

O        NO         t^»        ^         O 

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NO 

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CN1 

to 

»s§ 

to          W           ^"          ^"          CO 

* 

10 

Tj" 

* 

=|||<g 

M              t^            Tj-            CO          NO 
•*           -«4-           ON          CO           M 

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ON 
00 

M 

«. 

CO 

M 

VE'S'sli 
G  C7*^  O  oj 

0       NO        r^       to       0 

ON 

NO 

t^ 

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M.2 

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O         O        to       O      QtoOtotoco 

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to 

«l     1 

M              M              0              M          M     0 

M     0 

—    — 

H     0 

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HI     a 

00           O          00           Tf          Tj- 

oo         cs         co        to        rj- 

O          00            ON          O            O 

CN 

<N 

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to 

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d       d       d       d       d 

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t^       NO          0         «         to 

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M              M              M              M              M 

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to 

00 

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Tj-             H              -^-            M              M 

M 

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to        O         to        cs         O 

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CM              CO            N              t^-            CO 

j^ 

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CO          M           CO          C<           M 

CM 

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OO            to         OO            O            to 

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(N            rj-           CS            ON          Tf 

M         O         *™*         O         O 

d       d       d      d      d 

d 

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t^ 

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**         p 

:      :      :      :     8 

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frl   C3                 fee 

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•                 •                 •                 •              M 

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M 

o          5 

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M 
M 

8 

3        & 

....           10 

10 

»o 

ON 

NO 

- 

11 


RR 
6d 
XX 

.§.§ 
II 

oo 


II 


322  METHODS  IN  CHEMICAL  ANALYSIS 

present  with  a  fair  degree  of  accuracy,  and  to  separate  and 
determine  copper  when  associated  with  both  arsenic  and  anti- 
mony. In  the  latter  case  the  sum  of  the  arsenic  and  antimony 
may  also  be  determined,  but  the  values  here  obtained  for  copper 
tend  to  come  a  little  too  high  and  those  for  arsenic  and  antimony 
a  little  too  low. 

The  Estimation  of  Arsenic,  Antimony  and    Tin  in   the  Lower 
Condition    of   Oxidation    by    the    Action    of   Potassium 
Ferricyanide  in  Alkaline  Solution  and  Potas- 
sium Permanganate  in  Acid  Solution. 

In  1892  Quincke  *  published  a  method  for  estimating  arsenic 
and  antimony  gasometrically,  consisting  essentially  in  oxidizing 
the  arsenic  or  the  antimony  by  a  known  excess  of  potassium 
ferricyanide  in  the  presence  of  alkali,  and  determining  the  excess 
by  measuring  in  a  gasometer  the  oxygen  evolved  by  the  action 
of  hydrogen  peroxide  on  it  according  to  the  following  equation : 

2  K3FeC6N6  +  H2O2  +  2  KOH  =  2  K4FeC6N6  +  2  H2O  +  O2. 

The  same  process  of  oxidation  by  use  of  the  ferricyanide  has 
been  utilized  by  Palmer  f  in  the  estimation  of  arsenic,  antimony 
and  tin;  but,  to  determine  the  amount  of  ferricyanide  changed 
to  ferrocyanide  in  the  operation,  use  is  made  of  the  process  of 
titration  in  acid  solution  by  potassium  permanganate  previously 
employed  by  Browning  and  Palmer  J  in  the  estimation  of  cerium 
and  thallium.  In  alkaline  solution  the  oxidations  of  the  arse- 
nious,  antimonous  and  stannous  salts  take  place  according  to  the 
following  equations: 

As203  +  4  K3FeC6N6  +  4  KOH  =  As2O5  +  4  K4FeC6N6+2  H2O 
Sb2O3  +  4  K3FeC6N6  +  4  KOH  =  Sb2O5  +  4  K4FeC6N6+2  H2O 
SnO  +  2  K3FeC6N6  +  2  KOH  =  SnO2  +  2  K4FeC6N6+  H2O. 

The  ferrocyanide  formed  is  oxidized  in  acid  solution  by  perman- 
ganate according  to  the  equation 

10  K4FeC6N6  +  2  KMnO4  +  8  H2SO4  = 

10  K3FeC6N6  +  6  K2SO4  +  2  MnSO4+8  H2O. 

*  Zeit.  anal.  Chem.,  xxxi,  I. 

t  Howard  E.  Palmer,  Am.  Jour.  Sci.  [4],  xxix,  399. 

j  See  pages  223,  249. 


ARSENIC,   ANTIMONY  AND   TIN 


323 


Determination  of  To  the  solution  containing  antimonious  chloride 
Antimonious1  dissolved  in  enough  hydrochloric  acid  to  prevent  the 
Condition.  formation  of  a  basic  salt,  are  added  at  least  five 
times  as  much  potassium  ferricyanide  as  is  theoretically  neces- 
sary and  about  25  cm.3  of  a  20  per  cent  solution  of  potassium 
hydroxide.  After  standing  a  few  minutes,  the  solution  is 
strongly  acidified  with  dilute  sulphuric  acid  and  titrated  with 
permanganate  without  previous  removal  of  the  antimonic  salt. 
Results  are  given  in  the  table. 

The  Antimonious  Salt. 


Sb2O3  taken, 
grm. 

K3FeC6N8. 
grm. 

KOH. 
grm. 

Volume. 
cm.3 

Sb2O3  found, 
grm. 

Error, 
grm. 

o  .  0986 

8 

4 

IOO 

O  .  0989 

+o  .  0003 

o  .  0986 

4 

4 

75 

o  .  0984 

—  O.OOO2 

o  .  0986 

2 

4 

75 

o  .  0984 

—  O.OOO2 

o  .  0986 

4 

4 

150 

o  .  0984 

—  O.OOO2 

0.0986 

4 

4 

75 

o  .  0984 

—  O.OOO2 

0.0493 

4 

4 

75 

o  .  0495 

+O.OOO2 

0.0493 

4 

4 

75 

0.0497 

+o  .  0004 

0.0493 

4 

4 

75 

0.0495 

+O.OOO2 

0.1479 

4 

4 

75 

o.  1482 

+  0.0003 

0.1479 

4 

4 

75 

0.1477 

—  O.OO02 

0.1479 

4 

4 

75 

o.  1476 

—  0.0003 

0.1971 

8 

8 

125 

0.1972 

+O.OOOI 

Determination  of     ^°  tne  solution  of  stannous  chloride  free  from 

Tin  in  stannous  oxygen  and  kept  under  an  atmosphere  of  hydrogen 

are  added  in  solution  at  least  five  times  as  much 

potassium  ferricyanide  as  is  theoretically  necessary  and  enough 

The  Stannous  Salt. 


Sn  taken, 
grm. 

KjFeC6N6. 
grm. 

KOH. 
grm. 

Volume. 
cm.3 

Sn  found, 
grm. 

Error, 
grm. 

0.1032 

2-5 

6 

65 

0.1033 

+O.OOOI 

O.  IO22 

2-5 

6 

65 

o.  1016 

—  0.0006 

o.  1029 

3 

7 

60 

o.  1030 

+O.OOOI 

o.  1009 

5 

6 

85 

o.  ion 

+  O.OO02 

0.1005 

5 

5 

60 

O.  IOIO 

+0.0005 

O.IOII 

10 

5 

85 

0.1015 

+O.O004 

0.0995 

IO 

5 

85 

o.  1004 

+O.OOO9 

O.2O2O 

4 

6 

80 

o.  2019 

—  O.OOOI 

0.2003 

5 

6 

90 

o.  1998 

-0.0005 

O.2O2I 

IO 

5 

85 

o.  2027 

+0.0006 

324  METHODS  IN  CHEMICAL  ANALYSIS 

of  a  solution  of  potassium  hydroxide  to  completely  dissolve  the 
precipitated  stannic  acid,  the  two  solutions  of  ferricyanide  and 
potassium  hydroxide  having  been  previously  mixed.  The  stannic 
salt  is  precipitated  by  the  addition  of  about  10  grm.  of  ammonium 
sulphate,  and  warming  to  50°  or  60°.  After  settling,  the  precip- 
itate is  filtered  off  on  asbestos,  under  gentle  pressure,  and  washed 
with  a  10  per  cent  solution  of  ammonium  sulphate.  The  filtrate  is 
strongly  acidified  with  sulphuric  acid,  and  titrated  with  perman- 
ganate. 

The  test  results  are  given  in  the  preceding  table. 
Determination         The  essential  procedure  in  estimating  arsenic  taken 
*n  tne  arsenious  form  is  to  oxidize  by  potassium 


Condition.  ferricyanide  in  presence  of  a  fixed  alkali  hydroxide, 

neutralize  the  last  by  the  addition  of  ammonium  sulphate,  precip- 
itate the  arsenic  by  magnesia  mixture,  filter  off  the  ammonium 
magnesium  arsenate,  and  titrate  the  filtrate  with  permanganate, 
after  acidification  with  sulphuric  acid. 

The  procedure  recommended  is  as  follows:  To  the  solution 
containing  the  arsenic  in  the  arsenious  condition  is  added  an 
amount  of  potassium  ferricyanide  equal  to  at  least  five  times  the 
amount  theoretically  required  to  oxidize  the  arsenic  to  the  higher 
condition  of  oxidation,  and  about  25  cm.3  of  a  20  per  cent  solu- 
tion of  potassium  hydroxide,  keeping  the  volume  of  the  solution 
less  than  100  cm.3  After  standing  a  few  minutes,  the  free  fixed 
alkali  is  removed  and  the  solution  made  ammoniacal  by  dis- 
solving in  it  about  10  grm.  of  ammonium  sulphate,  and  magne- 
sia mixture,*  about  100  cm.3,  is  added. 

After  settling,  the  ammonium  magnesium  arsenate  is  filtered 
off  on  asbestos,  and  washed  with  faintly  ammoniacal  water. 
The  filtrate  is  strongly  acidified  with  dilute  sulphuric  acid,  and 
titrated  with  permanganate. 

As  Griitzner  f  has  shown,  during  the  titration  of  large  amounts 
of  ferrocyanide  by  permanganate,  a  precipitate  of  K2MnFeC6N6 
often  forms  by  the  action  of  unchanged  ferrocyanide  on  the 
manganese  sulphate  reduced  from  the  permanganate.  This 
precipitate  slowly  clears  up  as  more  permanganate  is  added, 
disappearing  entirely  as  the  end  point  is  reached,  but  it  tends 

*  About  100  cm.3  of  the  mixture  made  up  of  MgCl  .6  H2O,  55  grm.;  NH4C1, 
29  grm.;   NH4OH,  5  cm.3,  in  a  liter  of  water. 
t  Chem.  Centralblatt,  1902,  i,  500. 


VANADIUM 


325 


to  cause  high  results  by  obscuring  the  end  point.  This  difficulty 
is  overcome,  however,  by  titrating  in  the  presence  of  a  large 
amount  of  sulphuric  acid,  the  formation  of  this  precipitate  being 
thus  prevented.  The  titration  may  safely  be  made  in  the  cold 
in  the  presence  of  10  per  cent  of  sulphuric  acid,  and  this  amount 
is  generally  sufficient  to  prevent  the  formation  of  the  precipitate. 
Experimental  results  are  given  in  the  table. 

Arsenious  Oxide. 


As2O3  taken, 
grm. 

K3FeC6N6. 
grm. 

KOH. 
grm. 

Volume. 
cm.3 

As2Os  found, 
grm. 

Error, 
grm. 

0.0499 

8 

•25 

100 

0.0502 

+0.0003 

0  .  0499 

4 

•25 

75 

0.0499 

O.OOOO 

o  .  0499 

3 

•25 

75 

0.0501 

+O  .  OOO2 

0.0501 

4 

•25 

75 

0.0500 

—  O   OOOI 

0.0997 

8 

•25 

IOO 

o  .  0999 

+O.OOO2 

0.0997 

6 

•25 

90 

0.0993 

—  o  .  0004 

0.0997 

6 

1-25 

90 

o  .  0993 

—  0.0004 

O.IOOI 

8 

1-25 

IOO 

o  .  0998 

—  O.OO03 

o.  1496 

9 

1-25 

no 

0.1492 

—  O.OOO4 

o.  1496 

10 

1-25 

no 

o  .  1496 

O.OOOO 

0.1502 

10 

1-25 

no 

o.  1502 

O.OOOO 

0.1994 

12 

1-25 

125 

0.1994 

O.OOOO 

0.  IOOI 

4 

4- 

75 

o  .  0998 

—0.0003 

O.IOOI 

3 

4- 

75 

o  .  0998 

—0.0003 

The  Estimation  of  Arsenic  Acid  and  Antimonic  Acid  Associated 
with  Vanadic  Acid. 

t  Methods  for  the  estimation  of  arsenic  acid  and  of  antimonic 
acid  in  association  with  vanadic  acid,  dependent  upon  differen- 
tial reduction  followed  by  reoxidation,  are  discussed  in  connec- 
tion with  methods  for  the  determination  of  vanadium.* 


VANADIUM. 

The  Gravimetric  Estimation  of  Vanadic  Acid  Based  on  Liberation 
of  Iodine  and  Absorption  of  that  Element  by  Silver. 

Perkins  has  shown  that  when  a  soluble  vanadate  is  added  to 
an  excess  of  potassium  iodide  made  acid  with  hydrochloric  and 
shaken  with  electrolytic  silver  in  an  atmosphere  of  hydrogen,  f 
iodine  is  set  free  and  then  is  absorbed  by  the  silver.  %  From  the 

*  See  page  350. 

t  Claude  C.  Perkins,  Am.  Jour.  Sci.  [4],  xxix,  540. 

t  See  Fig.  10,  pages  27,  444. 


326 


METHODS  IN  CHEMICAL  ANALYSIS 


increase  in  weight  of  the  silver  taken  the  amount  of  vanadium 
pentoxide  may  be  calculated  upon  the  assumption  that  one  mole- 
cule sets  free  two  atoms  of  iodine  according  to  the  equation 

V2O5  +2  KI  +  2  HC1  =  V2O4  +  H2O  +  2  KC1  +  I2. 

Results  obtained  in  this  manner  with  ammonium  vanadate, 
the  composition  of  which  had  been  determined  by  fusion  with 
slightly  acidic  sodium  tungstate,  are  shown  in  the  following 
table. 

Reduction  of  V anodic  Acid,  by  Hydriodic  Acid:   Absorption  of  Liberated  Iodine 

by  Silver. 


Ag  taken, 
gnu. 

V2O6  taken, 
grm. 

I  found, 
grm. 

Calculated  V2O6. 
grm. 

Error, 
grm. 

2.OOI2 

0.1946 

0.2699 

0.1939 

—  0  .  0007 

2.OOI2 

0.2051 

0.2856 

0.2051 

0.0000 

2.0024 

0.1964 

0.2746 

0.1972 

+0.0008 

2.0024 

0.2362 

0-3295 

0.2367 

+0.0005 

2.0062 

0-2953 

0.4106  . 

0.2949 

—  0.0004 

2.0062 

0.2681 

0.3729 

0.2679 

—  O.OOO2 

2.0527 

0.5770 

0.8038 

0-5774 

+  0.0004 

2.OOO6 

0.2035 

0.2829 

0.2032 

—  0.0003 

The  Precipitation  of  Ammonium  Vanadate  by  Ammonium 
Chloride. 

Berzelius  was  the  first  to  point  out  and  utilize  in  analysis  the 
fact  that  when  a  vanadate  in  concentrated  solution  is  treated 
with  a  saturated  solution  of  ammonium  chloride,  white  insoluble 
ammonium  meta vanadate  is  deposited.*  Modifications  of  the 
method  have  been  proposed  by  V.  Hauer,f  Ditte,t  and  Gibbs,§ 
all  of  which  have  been  unfavorably  criticized. || 

Gooch  and  Gilbert**  have  shown,  however,  that  the  process  of 
Gibbs  gives  a  practically  complete  precipitation  of  ammonium 
vanadate,  when  to  the  slightly  ammoniacal  solution  of  the  solu- 
ble vanadate  such  an  excess  of  ammonium  chloride  is  added  that 

*  Ann.  Phys.,  xcviii,  54,  1831. 
t  Jour,  prakt.  Chem.,  Ixix,  388. 
J  Compt.  rend.,  civ,  982. 
§  Proc.  Am.  Acad.,  x,  242,  249. 

||  Milch,  Inaug.    Diss.,   Berlin,   1887.     Rosenheim,   Inaug.   Diss.,   Berlin, 
1888.     Liebert,  Inaug.  Diss.,  Halle,  1891.     Euler,  Inaug.  Diss.,  Berlin,  1895. 
**  F.  A.  Gooch  and  R.  D.  Gilbert,  Am.  Jour.  Sci.  [4],  xiv,  205. 


VANADIUM  327 

the  solution  deposits  ammonium  chloride  after  concentration 
and  cooling,  and  the  mixture  is  then  allowed  to  stand  twenty- 
four  hours.  Should  too  much  ammonium  chloride  for  convenient 
handling  crystallize  out  on  cooling,  this  is  to  be  redissolved  by 
the  careful  addition  of  ammonia;  but  care  should  be  taken  that, 
after  standing,  a  little  solid  ammonium  chloride  and  free  ammonia 
still  remain  in  the  mixture.  The  precipitated  ammonium  vana- 
date  is  washed  with  a  cold  saturated  solution  of  pure  ammo- 
nium chloride,  and  the  vanadium  in  the  vanadate  determined 
by  any  appropriate  means.  Volumetric  processes  are  to  be 
preferred. 

According  to  this  procedure  the  vanadate  is  dissolved  by 
digestion  upon  the  steam  bath  in  25  cm.3  of  slightly  ammonia- 
cal  water.  The  mixture  is  allowed  to  remain  upon  the  steam 
bath,  with  addition  from  time  to  time  of  a  little  ammonia,  to 
keep  the  meta vanadate  of  normal  composition  and  colorless, 
until  the  volume  has  been  reduced  to  about  25  cm.3.  On  cooling, 
a  small  amount  of  ammonium  chloride  crystallizes  out,  but  only 
a  little  if  the  proportion  has  been  properly  adjusted.  If  too 
large  an  amount  crystallizes  out,  it  is  nearly  redissolved  by  the 
cautious  addition  of  ammonium  hydroxide.  The  mixture  is 
allowed  to  stand  twenty-four  hours  to  insure  complete  crystal- 
lization of  the  ammonium  metavanadate,  and  is  then  filtered  on  a 
weighed  asbestos  filter  in  the  perforated  crucible,  the  precipitate 
being  transferred  and  washed  with  a  cold  saturated  solution  of 
ammonium  chloride.  Crucible  and  precipitate  are  heated,  at 
first  very  gently  to  drive  off  the  ammonium  chloride  without 
occasioning  mechanical  loss  of  the  vanadium,  and  finally  to  red- 
ness and  fusion  of  the  pentoxide  remaining. 

Without  special  precaution  some  trouble  is  occasionally  found 
in  removing  from  the  walls  of  the  beaker  the  adherent  crystals 
of  ammonium  vanadate,  but  the  difficulty  may  be  overcome 
by  the  device  of  forming  upon  the  walls  of  the  beaker  before 
using  it  a  film  of  paraffin  of  extreme  thinness  by  rinsing  the 
beaker  with  a  dilute  solution  of  paraffin  in  naphtha  (0.5  grm.  of 
paraffin  in  300  c.c.  of  naphtha)  and  allowing  the  naphtha  to 
evaporate.  Crystals  of  the  vanadate  adhering  to  the  walls  of 
the  beaker  thus  previously  prepared  are  easily  removed  by 
means  of  the  ordinary  rubber  or  "policeman."  The  table  con- 
tains the  record  of  six  consecutive  experiments  made  after  some 


328 


METHODS  IN   CHEMICAL  ANALYSIS 


preliminary  study  of  the  method.  Tests  of  the  washings  and 
nitrate  were  made  in  several  instances  by  acidifying  with  hydro- 
chloric acid  and  testing  with  hydrogen  dioxide,  but  no  indication 
of  the  presence  of  vanadium  was  obtained. 

Precipitation  by  Ammonium  Chloride  and  Ammonia:  Gravimetric  Determination. 


NH4VO8  taken.* 
grm. 

V2O5  present, 
grm. 

V2O5  found, 
grm. 

Error, 
grm. 

0-5 

0.3807 

0.3814 

+0.0007 

0-5 

0.3807 

0.3818 

-f-O.OOII 

o-S 

0.3807 

0-3813 

+0.0006 

0-5 

0.3807 

0.3808 

+O.OOOI 

0-5 

0.3807 

0.3808 

+O.OOOI 

o-S 

0.3807 

0.3799 

—0.0008 

*  Containing  76.14%  of  V2O6,  according  to  Holverscheit's  method. 

These  results  are  sufficient  to  show  that  the  method  of  Gibbs  is 
capable  of  yielding  an  analytical  separation  of  value,  but,  as  Gibbs 
pointed  out,  it  is  ordinarily  preferable  to  estimate  the  vanadium 
volumetrically  rather  than  to  go  through  the  tedious  and  exact- 
ing process  of  ignition  to  recover  the  vanadium  pentoxide. 

Other  results  obtained  in  determining  by  the  Holverscheit 
process  the  ammonium  vanadate  precipitated  by  the  Gibbs  proc- 
ess confirm  the  gravimetric  results. 

Precipitation  by  Ammonium  Chloride  and  Ammonia:    Volumetric 
Determination. 


NH4VO3  taken, 
grm. 

V2O5  found. 

grin. 

Error, 
grm. 

O.  1000 

0.0760 

—  O.OOOI 

O.  1000 

0.0764 

+0.0003 

O.  IOOO 

0.0763 

+  O.O002 

O.  IOOO 

0.0962 

+O.OOOI 

O.  IOOO 

0.0757 

—  O.OOO4 

O.  IOOO 

0.0751 

—  O.OOIO 

O.  IOOO 

0.0761 

0  .  OOOO 

O.  IOOO 

0.0758 

—0.0003 

The  Estimation  of  Vanadium  as  Silver  Vanadate. 

The  conditions  under  which  vanadium  may  be  estimated  as 
silver  vanadate  have  been  investigated  by  Browning  and  Palmer.* 
The  experimental  results  indicate  that  the  composition  of  the 

*  Philip  E.  Browning  and  Howard  E.  Palmer,  Am.  Jour.  Sci.  [4],  xxx,  220. 


VANADIUM 


329 


precipitate  thrown  down  in  solutions  acidified  with  acetic  acid 
is  variable,  while  that  of  the  precipitate  from  neutral  solution 
has  after  ignition  the  composition  of  the  metavanadate  AgVOa. 

The  preferred  procedure  is  as  follows:  The  solution  of  the 
vanadate  is  made  acid  with  nitric  acid  and  then  alkaline  with 
ammonia.  It  is  boiled  until,  when  the  ammonia  is  almost 
expelled,  the  solution  begins  to  turn  yellow.  At  this  point  the 
boiling  is  stopped  because  if  it  .is  continued  further  the  solution 
becomes  too  acidic  and  the  precipitate  which  forms  on  addition 
of  silver  nitrate  is  no  longer  definitely  the  metavanadate.  To  the 
solution,  about  200  cm.3  in  volume  and  heated  to  boiling,  a  solu- 
tion of  silver  nitrate  is  added  with  vigorous  stirring  to  coagulate 
the  precipitate.  The  precipitate  is  then  filtered  off  on  an  as- 
bestos  felt  contained  in  a  perforated  crucible,  washed  thoroughly, 
ignited  at  a  gentle  heat  below  the  fusing  point  of  silver  vanadate, 
and  weighed  as  AgVO3. 

Following  are  results  of  this  procedure : 

Precipitation  as  Silver  Metavanadate. 


V2O5  taken. 

AgVO3  found. 

V2O5  found. 

Error. 

grm. 

grin. 

grm. 

gnu. 

Precipitated  from  neutral  solution  by  AgNOs. 

0.1139 
0.0569 
0.0569 
0.0569 
0.1066 

0-2595 
0.1291 
0.1277 
0.1303 
0.2436 

0.1143 
0.0569 
0.0562 
0.0574 
0.1073 

+O.OOO4 
+O.OOOO 

+0.0007 

+0.0005 
+0.0007 

Ammoniacal  solution  boiled  to  expel  ammonia:   AgNO3  added  to  boiling 

solution. 


0.1066 

0.2430 

0.1070 

+o  .  0004 

0.1066 

0.2429 

0.1070 

+0.0004 

0.0533 

0.1224 

0.0539 

+0.0006 

0.0533 

O.I22I 

0.0538 

+0.0005 

0.0569 

0.1312 

0.0578 

+0.0009 

Solution  acid  with  HNO3  made  ammoniacal  and  boiled  to  expel  ammonia: 
AgNO3  added  as  soon  as  the  solution  began  to  turn  yellow. 


0.0569 

0.1293 

0.0569 

o.oooo 

0.1066 

0.2419 

0.1065 

—  o.oooi 

0.1066 

0.2424 

0.1068 

+0.0002 

0.1066 

0.2422 

0.1066 

+O.OOOO 

0.0533 

0.1215 

0.0535 

+O.OOO2 

330  METHODS  IN  CHEMICAL  ANALYSIS 

The  Estimation  of  V r anodic  Acid  by  the  Action  of  the  Halogen 

Acids. 

The  Action  of  The  estimation  of  vanadic  acid  by  reduction  with 
Hydrochloric  concentrated  hydrochloric  acid  followed  by  the 
iodometric  determination  of  the  chlorine  evolved, 
as  proposed  by  Bunsen  *  and  noted  by  Mohr,f  has  been  unfavor- 
ably criticized  by  many  investigators.!  On  the  other  hand, 
Gibbs  §  has  determined  the  small  amounts  of  vanadium  pen- 
toxide  found  in  the  vanadio-tungstates  and  other  complex  com- 
binations, by  boiling  with  strong  hydrochloric  acid,  collecting 
in  potassium  iodide  the  chlorine  evolved,  titrating  the  freed 
iodine  with  sodium  thiosulphate  and  calculating  from  the  amount 
of  iodine  thus  found  the  vanadium  pentoxide  corresponding  to  a 
change  of  condition  from  V2O5  to  V2O4. 

Gooch  and  Stookey  ||  have  shown  that  when  the  action  of 
hydrochloric  acid  is  sufficiently  continued,  vanadic  acid  may  be 
completely  reduced  to  the  condition  of  the  tetroxide ;  and  have 
pointed  out  that  the  method  of  reduction  is  of  special  advantage 
in  those  cases  which  call  for  titration  of  the  tetroxide  by  per- 
manganate, since  in  such  cases  the  use  of  Holverscheit's  admi- 
rable method  of  reduction  by  hydrobromic  acid  is  precluded.  It 
appears  also  that  it  is  possible  to  effect  the  reduction  of  vanadic 
acid  to  within  a  few  per  cent  of  the  amount  present  by  a  single 
treatment  with  concentrated  hydrochloric  acid,  and  that  the 
amount  of  the  reduction  may  be  determined  by  titrating  the 
residue  with  potassium  permanganate. 

In  the  action  of  the  most  highly  concentrated  aqueous  hydro- 
chloric acid  upon  vanadic  acid  three  stages  are  marked :  First,  the 
vanadic  acid  dissolves  to  a  solution  of  a  deep  brown  color  with- 
out perceptible  evolution  of  chlorine;  upon  warming,  the  solution 
evolves  chlorine  and  takes  on  a  deep  green  color;  thereafter  the 
evolution  of  chlorine  becomes  weaker,  and  the  solution,  giving 
off  a  small  amount  of  chlorine  on  strong  boiling,  assumes  a  clear 

*  Ann.  Chem.,  xcvi,  265. 

f  Titrirmethode,  vte  Aufl.,  314. 

J  Czudnowicz,  Ann.  Phys.,cxx,  17.  Milch,  Inaug.  Diss.,  Berlin,  1887,  10. 
Rosenheim,  Inaug.  Diss.,  Berlin,  1888;  Ann.  Chem.,  ccli,  197.  Holverscheit, 
Inaug.  Diss.,  Berlin,  1890. 

§  Proc.  Am.  Acad.,  x,  250. 

||  F.  A.  Gooch  and  L.  B.  Stookey,  Am.  Jour.  Sci.,  [4],  xiv,  369. 


VANADIUM  331 

blue  color.  Observing,  however,  that  the  liberation  of  chlorine 
actually  begins  to  some  extent  as  soon  as  the  hydrochloric  acid 
and  vanadic  acid  come  into  contact,  Gooch  and  Stqokey  have 
used  the  form  of  apparatus  shown  in  Fig.  3,  p.  4,  so  that  no 
chlorine  may  be  evolved  before  the  apparatus  is  connected  and 
ready.  For  the  retort  is  used  a  Voit  flask  upon  the  inlet  tube  of 
which  is  sealed  a  stoppered  funnel,  while  the  outlet  tube  is  sealed 
to  a  Drexel  wash  bottle  used  as  a  receiver,  and  this  in  turn  is 
joined  to  Will  and  Varrentrapp  bulbs.  Connection  is  provided 
between  the  funnel  and  a  carbon  dioxide  generator,  so  that  by 
passing  carbon  dioxide  through  the  apparatus  all  danger  of 
regurgitation  may  be  avoided.  A  branched  connection  with  the 
funnel  tube,  so  arranged  that  either  hydrochloric  acid  gas  or 
carbon  dioxide  or  both  may  at  pleasure  be  sent  into  the  apparatus, 
is  a  convenience. 

According  to  the  procedure  described,  the  receiver  and  trap 
are  charged  with  water  containing  3  grm.  of  potassium  iodide. 
The  vanadate  is  introduced  into  the  dry  flask  and  the  apparatus 
ii  adjusted.  Hydrochloric  acid  is  put  in  the  stoppered  separating 
funnel  and,  when  all  is  ready,  allowed  to  run  into  the  flask  and 
boiled  until  the  liquid  takes  a  .  clear  blue  color.  Thereafter, 
either  the  evaporation  is  continued  to  dryness  and  fresh  con- 
centrated hydrochloric  acid  added,  or  else  the  weak  acid  remain- 
ing after  the  boiling  process  is  cooled  and  resaturated  with 
gaseous  hydrochloric  acid;  and,  in  either  case,  the  boiling  is  re- 
peated. 

Several  successive  treatments  with  strongest  hydrochloric  acid, 
consisting  either  in  the  addition  of  the  concentrated  aqueous  acid 
to  the  dry  residue  in  successive  portions  or  in  passing  gaseous 
acid  through  the  residual  liquid  cooled  in  ice  water,  finally  bring 
the  residue  to  a  condition  of  reduction  in  which  the  recharging 
of  the  liquid  residue  with  gaesous  acid  produces  no  brown  or 
green  color,  but  a  clear  blue.  When  this  condition  is  reached  the 
reduction  is  complete,  the  application  of  the  Holverscheit  pro- 
cedure to  the  residue  producing  no  appreciable  further  reduc- 
tion. The  holding  of  the  blue  color,  when  the  solution  cooled 
in  ice  water  is  thoroughly  charged  with  the  gaseous  acid,  appears 
to  be  the  best  evidence  of  the  complete  reduction  of  V2O5  to 
V2O4;  it  is  not  sufficient  that  the  boiled  solution,  in  which  the 
acid  has  been  weakened,  should  be  blue.  The  determination  of 


332 


METHODS  IN  CHEMICAL  ANALYSIS 


o 

=8 

^ 

•e 

o 

1 

^ 


>l 


J 


tf-g 


. 

W    W 

w   <D 


•  +J 
G*rt 


\      M  W 

i     •  ,Q  ctf 

I  o  <u  1-1 

;1  '8 .« 


WWW 


pqpqpq 


:8     88 


000 

I  I  I 


o       o 
d      6 


<*J* 

I 


fO  10 
O   O 

odd 


poo 
odd 


8     8 


ass    a    a  a  a 


o  o  o 

#    # 
O  »o  to 


0    CO    W 

goo 


VANADIUM 


333 


the  free  iodine  in  the  receiver  by  titration  with  thiosulphate 
measures  the  amount  of  vanadate  present,  according  to  the 
equations 

V205  +  2  HC1  =  V204  +  H20  +  C12 
C12  +  2  KI     =2  KC1  +  I2. 

Test  experiments  made  with  ammonium  vanadate  analyzed  by 
the  Holverscheit*  process,  are  given  in  the  preceding  table. 

The  results  of  other  similar  experiments  in  which  the  reduced 
vanadium  tetroxide  in  the  residue  was  determined  by  titration 
with  potassium  permanganate  after  the  addition  of  a  manganous 
salt  are  given  in  the  following  table ;  and  in  some  cases  the  results 
of  the  iodometric  determination  of  chlorine  set  free  are  brought 
into  comparison  with  the  permanganate  determination  of  the 
tetroxide. 

Action  of  Hydrochloric  Acid  upon  the  Vanadate:    Oxidation  of  the  Residue  by 

Permanganate. 


NH4V03. 

grin. 

HC1. 
cm.8 

V206 

calculated, 
grin. 

V2O5  found 
by  chlorine 
in  distillate. 

grm. 

Error, 
grm. 

V2O«;  found 
by  KMn04 
as  residue. 

grm. 

Error, 
grm. 

O    IOOO 

2S,  sp.  gr.  i  20 

0.0765 

o  0740* 

—  o  0025 

O    IOOO 

25,  sp.  gr.  i  .  20 

0.0765 

o  0736* 

—  O   OO2Q 

O    IOOO 

2$,  sp.  gr.  i  .  20 

0.0765 

o  0744* 

—  O   OO2I 

O.IOOO 

30,  sp.  gr.  i  .  20 

0.0765 

0.0711 

—  0.0054 

* 

Gas.t 

O  O734 

—  o  003  1 

O   O7^8t 

—  o  0007 

O.  IOOO 

25,  sp.  gr.  i  .20 

0.0765 

0.0714 

—  O.OO5I 

* 

Gas.t 

0.0746 

—  o  0019 

* 

Gas.t 

O.O761? 

o  oooo 

O  076^1" 

o  oooo 

0.3000 

2<;,  sp.  srr.  i  .  20 

o.  2201; 

* 

Gas  i 

* 

o.  3000 

Gas.t 
25,  sp.  gr.  i  .  20 

O.229<J 

O.  22^Q 

—  0.0036 

0.2284! 

* 

—  0.0011 

Gas.t 

0.2288 

—  0.0007 

o  2302! 

~f"O   OOO7 

0.3000 

25,  sp.  gr.  1.17 
Gas.t 

0.22Q5 

0.2207 

o  224.4. 

-0.0088 

—  o  0051 

* 
* 

Gas.t 

o  2258 

—  o  .  003  7 

* 

Gas.t 

o  2269 

—  o  0026 

* 

Gas.t 

o  2281 

—  o  0014 

O    22Q^t 

—  O   OOO2 

*  Residue  after  boiling  was  blue. 

t  Residue  blue  alter  boiling  was  still  blue  when  resaturated. 

t  Residue  as  cooled  in  ice  water  and  re-saturated  with  gaseous  HC1. 

*  Inaug.  Diss.,  Berlin,  1890,  p.  48. 


334  METHODS   IN   CHEMICAL  ANALYSIS 

When  hydrochloric  acid  of  suitable  concentration  and  the  van- 
adate  come  into  contact,  the  liberation  of  chlorine  is  immediate, 
some  of  which  escapes  from  the  solution  while  some  is  retained, 
and  the  reaction  proceeds  to  a  balance  as  indicated  in  the  ex- 
pression 

V2O5  +  2  HC1  <=»  V2O4  +  H2O  +  C12. 

To  complete  the  reduction  of  the  higher  oxido  it  is  necessary 
to  remove  the  free  chlorine  from  the  system  while  keeping  up  the 
requisite  strength  of  the  hydrochloric  acid.  In  removing  the 
chlorine  by  boiling,  the  concentration  of  the  hydrochloric  acid 
is  diminished  below  the  point  at  which  action  upon  vanadic  acid 
may  take  place  with  liberation  of  chlorine.  This  h  why,  in  push- 
ing the  action  to  completion  by  the  boiling  process,  it  is  neces- 
sary to  increase  the  concentration  of  the  hydrochloric  acid  from 
time  to  time,  either  by  cooling  and  recharging  with  gaseous  acid 
or  by  evaporating  off  the  weak  acid  and  replacing  it  by  strong 
acid. 

In  continuing  the  study  of  this  reaction,  Gooch  and  Curtis* 
have  tried  to  complete  the  reaction  by  bubbling  a  current  of 
gaseous  hydrochloric  acid  through  the  cooled  residue  of  a  single 
treatment  by  boiling.  Under  these  conditions  the  hydrochloric 
acid  is  always  at  the  concentration  of  activity,  but  the  re- 
moval of  the  chlorine  is  necessarily  slow  since  the  current  of  gas 
must  not  be  rapid  enough  to  cause  mechanical  loos  from  the 
mixture. 

The  apparatus  used  was  that  described  above  and  the  vanadate 
was  standardized  by  the  method  of  Ilolverscheit. 

In  every  experiment  approximate1^  o.i  grm.  of  ammonium 
vanadate  was  first  introduced  into  the  reduction  flask  B.  The 
air  was  expelled  from  the  apparatus  by  carbon  dioxide  from  the 
generator,  the  receiver  C  being  charged  with  hydrochloric  acid 
and  the  trap  g  with  water.  Concentrated  hydrochloric  acid 
(15  cm.3)  was  admitted  through  the  stoppered  funnel  A,  and  the 
mixture  was  boiled.  The  deep  red  color,  produced  when  the  acid 
was  first  added,  gradually  passed  through  green  to  blue.  The 
flask  was  allowed  to  cool,  carbon  dioxide  being  admitted  to  fill 
the  partial  vacuum,  and  surrounded  with  ice.  Hydrochloric 
acid  gas  was  passed  through  the  solution  in  the  reduction  flask, 

*  F.  A.  Gooch  and  R.  W.  Curtis,  Am.  Jour.  Sci.,  [4],  xvii,  41. 


VANADIUM 


335 


at  the  rate  of  one  or  two  bubbles  a  second,  for  periods  varying 
from  J  hour  to  H2j  hours,  the  solution  turning  brown  at  first 
and  then  changing  to  green  or  blue,  according  to  the  length  of 
the  period.  For  continuing  the  flow  of  gas  for  long  periods 
small  Kipp  generators  set  up  with  massive  ammonium  chloride 
and  concentrated  sulphuric  acid  were  found  very  convenient, 
a  single  charge  serving  to  keep  up  the  flow  continuously  over 
night. 

At  the  end  of  the  operation  the  degree  of  reduction  was  deter- 
mined by  titrating  the  contents  of  the  flask,  after  dilution,  with 
potassium  permanganate  in  presence  of  a  man^anous  salt.  The 
data  of  the  experiments  are  detailed  in  the  following  table : 

Reduction  by  Hydrochloric  Acid  at  Room  Temperature. 


V2O4  per  o.iooo  grm.  of 

NH4VO, 

taken. 

Time. 

vanadate  taken. 

Difference. 

Calculated. 

Found. 

gnu. 

hrs. 

O.IO22 

* 

0.0695 

0.0619 

0.0076 

O.II2I 

17 

0.0695 

0.0621 

0.0074 

0,1010 

i8| 

0.0695 

0.0678 

0.0017 

o  .  0980 

21 

0.0695 

0.0658 

0.0037 

o.  1044 

30 

o  .  0695 

o  .  0600 

0.0005 

0.1043 

112^ 

0.0605 

0.0691 

o  .  0004 

The  reduction  of  the  vanadic  acid  to  the  condition  of  tl.e 
tetroxide  by  the  action  of  hydrochloric  acid  in  the  cold  is  obvi- 
ously slow,  as  would  be  expected,  but  the  results  show  that  it 
may  be  made  practically  complete  in  this  manner.  No  indica- 
tion of  reduction  below  the  condition  of  the  tetroxide  by  the 
agency  of  hydrochloric  acid  is  apparent. 

In  Holverscheit's  method  the  reduction  of  vanadic 
acid  by  the  action  of  hydrochloric  acid  and  small 
amounts   of    potassium   bromide   to   the   condition 
represented  by  vanadium  tetroxide  is  ideally  complete.     Under 
these  conditions  the  concentration  of  hydrobromic  acid  is  low. 
The  effect  of   more  concentrated  hydrobromic  acid  upon  the 
course  of  reduction  has  been  investigated  by  Gooch  and  Curtis.* 
In  the  experiments  recorded  in  the  following  table,  weighed  por- 
tions of  ammonium  vanadate  were  introduced  into  the  reduction 
*  F.  A.  Gooch  and  R.  W.  Curtis,  Am.  Jour.  Sci.,  [4],  xvii,  43. 


The  Action  of 
Hydrobromic 
Acid. 


336 


METHODS  IN  CHEMICAL  ANALYSIS 


flask,  the  receiver  and  trap  were  charged  with  a  solution  of 
potassium  iodide  (3  grm.  :  350  cm.3),  the  apparatus  was  filled 
with  carbon  dioxide,  hydrobromic  acid  (15  cm.3)  of  sp.  gr.  1.68 
(made  by  distilling  a  mixture  of  potassium  bromide  and  sirupy 
phosphoric  acid)  was  introduced  through  the  funnel  and  the 
mixture  was  boiled  eight  or  ten  minutes.  On  the  addition  of 
the  acid  the  vanadate  dissolved  and  the  solution  took  on  a  light 
green  color,  which  on  heating  changed  to  red-brown  and  finally 
to  a  clear  deep  green.  After  cooling,  the  degree  to  which  the 
vanadic  acid  had  been  reduced  was  estimated  in  two  ways  — 
by  determining  by  means  of  standard  sodium  thiosulphate  the 
iodine  set  free  in  the  receiver  by  the  bromine  evolved,  and  by 
oxidizing  by  standard  iodine  the  reduced  product  in  the  flask. 
The  latter  process  followed  the  lines  recommended  by  Browning* 
and  consisted  essentially  in  neutralizing  the  acid  in  the  reduction 
flask  by  potassium  bicarbonate,  adding  an  excess  of  twentieth 
normal  iodine,  allowing  the  mixture  to  stand  twenty  minutes 
(all  without  admission  of  air),  then  transferring  to  a  larger  flask, 
introducing  a  slight  excess  of  twentieth  normal  arsenite,  and 
titrating  with  iodine  in  presence  of  starch. 

The  results  of  these  experiments  are  calculated  upon  the 
hypotheses  (first)  that  the  vanadic  acid  is  reduced  to  the  tetrox- 
ide,  (second)  that  it  is  reduced  to  the  trioxide,  and  (third)  that 
the  trioxide  and  tetroxide  are  left  in  mixture. 

Reduction  by  Hydrobromic  Acid. 


Ino.ioooNH4V03 

taken. 

Found;  calculated  as 
V204. 

Found;  calculated  as 
V203. 

Found;  calculated  as 
mixture  from  figures 
for  receiver. 

V20<. 

V20,. 

Flask. 

Receiver. 

Flask. 

Receiver. 

V204. 

V20S. 

grin. 

gnu. 

gmi. 

grm. 

grm. 

grm. 

grm. 

grin. 

0.0699 

0.0632 

0.0913 

0.0877 

0.0413 

O  .  0396 

0.0521 

0.0160 

0.0699 

0.0632 

0.0879 

0.0885 

0.0397 

O  .  0400 

0.0513 

0.0168 

0.0699 

0.0632 

o  .  0849 

0.0896 

0.0384 

o  .  0405 

0.0502 

0.0179 

o  .  0699 

0.0632 

0.0858 

o  .  0849 

0.0388 

0.0384 

0.0549 

0.0136 

o  .  0699 

0.0632 

0.0854 

0.0853 

0.0386 

0.0385 

0.0545 

0.0139 

0.0699 

0.0632 

0.0841 

o  .  0839 

0.0380 

0.0379 

0.0559 

0.0127 

In  other  experiments  the  aqueous  hydrobromic  acid  remaining 
after  boiling  was  strengthened  by  cooling  and  recharging  with  the 

*  See  page  345- 


VANADIUM 


337 


gaseous  acid  liberated  by  heating  strong  aqueous  hydrobromic 
acid,  and  then  reboiled.  The  bromine  evolved  in  each  boiling 
was  determined,  and  finally  the  degree  of  reduction  of  the 
residue  was  estimated. 

Reduction  by  Hydrobromic  Acid. 


In  o.iooo  NH4VO3 
taken. 

Found;  calculated  as 
V204. 

Found;  calculated  as 
V302. 

Found;  calculated  as 
mixture  from  figures 
for  receiver. 

V204. 
grm. 

V203, 
grm. 

Flask, 
grm. 

Receiver, 
grm. 

Flask, 
grm. 

Receiver, 
grm. 

V204. 

grm.. 

V203. 
giro. 

o  .  0699 

0.0632 

* 
0.0945 

* 
* 
* 

0.1258 

0.0426 

0.0389 
0.0499 
0.0584 
0.0569 

0.0445 

0.0538 
0.0295 
0.0107 
0.0107 

O.O22I 

O.OI46 
0.0366 
0-0535 
0.0535 

0.0943 

0.0860 
0.1104 
0.1291 
o.  1291 

0.0427 

0.0699 

0.0632 

The  Action  of 


*  Recharged  with  HBr. 

So  it  appears  that  increase  in  concentration  of  the  hydro- 
bromic acid  tends  to  carry  the  reduction  below  the  condition  of 
the  tetroxide.  The  highest  degree  of  reduction  reached  in  these 
experiments  corresponds  to  a  mixture  of  one-sixth  tetroxide  and 
five-sixths  trioxide. 

Browning  *  has  shown  that  in  accordance  with  the 
theory  of  reduction  given  in  the  following  equation, 

V205  +  2  HI  =  V204  +  H20  +  I2, 

good  analytical  results  are  obtained  when  solutions  of  the  vana- 
date,  I  or  2  grm.  of  potassium  iodide  and  10  cm.3  of  sulphuric 
acid  of  half  strength  are  boiled  to  a  volume  of  about  35  cm.3  and 
the  residual  solution  is  cooled,  neutralized  with  an  alkali  bicar- 
bonate (after  the  addition  of  a  tartrate  to  prevent  precipitation), 
and  treated  for  some  time  with  an  excess  of  iodine  which  is  fol- 
lowed by  an  excess  of  arsenious  acid,  the  last  being  titrated  by 
iodine  in  presence  of  starch. 

In  Browning's  process  the  estimation  of  the  reduced  product 
in  the  residue  is  made  the  measure  of  action.  The  determination 
of  the  iodine  set  free  by  the  action  of  hydriodic  acid  has  been 
studied  by  Gooch  and  Curtis,  f  The  experiments  recorded  be- 

*  See  page  343. 

f  F.  A.  Gooch  and  R.  W.  Curtis,  Am.  Jour.  Sci.,  [4],  xvii,  45. 


338 


METHODS  IN  CHEMICAL  ANALYSIS 


low  were  conducted,  in  the  apparatus  of  Figure  3,  in  an  atmosphere 
of  carbon  dioxide.  The  determinations  of  the  iodine  collected  in 
potassium  iodide  in  the  receiver  are  compared  with  determina- 
tions of  the  reduction  in  the  residue.  The  determinations  of 
reduced  vanadic  oxide  were  made  either  by  neutralization  with  po- 
tassium bicarbonate,  by  addition  of  a  measured  amount  of  stand- 
ard iodine  followed  by  a  measured  amount  of  standard  arsenite 
and  back  titration  with  iodine,  or  by  treatment  of  the  residue 
with  iodine  before  neutralizing  with  the  bicarbonate  (to  avoid 
atmospheric  oxidation),  addition  of  arsenite  and  final  titration 
with  iodine. 

Reduction  by  Sulphuric  Acid  and  Potassium  Iodide  in  Moderate  Concentrations. 


V204  in 
o.iooo  grm. 
IS1H4VO3 

KI. 

H2SO4 
i  :  i. 

Initial 
volume. 

Final 

volume. 

Reduction  flask  V2O4. 

Receiver  V2O4. 

taken. 

Found. 

Error. 

Found. 

Error. 

grm. 

grm. 

cm.* 

cm.s 

cm. 

grm. 

grm. 

grm. 

grm. 

The  residue  neutralized  before  addition  of  iodine. 


0.0699 

10 

35 

0.0668 

—  0.0031 

O.O7OO 

+0  .  OOOI 

o  .  0699 

IO 

. 

35 

0.0692 

—  0.0007 

0.07IS 

-}-o.ooi6 

0.0699 

IO 

35 

o  .  0686 

—0.0013 

O.O7I8 

+0.0019 

o  .  0699 

TO 

50 

35 

0.0696 

-0.0003 

0.0744 

-j-o  .  0045 

o  .  0699 

IO 

45 

35 

0.0678 

—  O.CO2I 

0.0704 

+0.0005 

0.0699 

10 

So 

35 

0.0690 

—  0.0009 

O.O7IO 

+O.OOII 

o  .  0699 

IO 

60 

35 

0.0681 

—  O.OOlS 

0.0738 

+o  .  0039 

o  .  0699 

IO 

55 

35 

0.0689 

—  o.ooio 

0.0724 

+0.0025 

o  .  0699 

0.6 

6 

55 

35 

0.0679 

—  O.OO20 

O.O722 

+0.0023 

Iodine  added  before  neutralizing  the  residue. 


0.0699 

I 

10 

55 

35 

c  .  0699 

o  .  oooo 

0.0713 

+0.0014 

o  .  0699 

I 

6 

55 

35 

0.0713 

+0.0014 

0.0722 

+0.0023 

o  0690 

I 

IO 

80 

"2  ^ 

O   O72S 

+  O  .  OO  2  6 

w 
o  .  0699 

I 

IO 

75 

GO 
35 

0.0710 

+O.OOII 

w  •  w  /  *0 

0.0718 

+0.0017 

o  .  0699 

0.6 

4 

55 

35 

0.0701 

+  O.OOO2 

0.0709 

+  O.OOIO 

o  .  0699 

0.6 

IO 

55 

35 

0.0717 

+0.0018 

0.0745 

+0.004.6 

o  .  0699 

0.6 

6 

55 

35 

0.0706 

+0.0007 

0.0734 

+0.0035 

o  .  0699 

0.6 

6 

55 

35 

0.0703 

+0.0004 

0.0727 

+0.0028 

o  .  0699 

06 

4 

55 

35 

0.0700 

+  0.0001 

0.0731 

+0.0052 

In  those  experiments  in  which  the  vanadate  was  treated  with 
dilute  sulphuric  acid  and  potassium  iodide,  the  iodine  found  in 
the  receiver  indicates  a  trifling  reduction  beyond  the  condition  of 
the  tetroxide  V2O4,  averaging  0.0023  grm.;  and  the  same  in 
general  is  true  in  regard  to  those  determinations  of  the  residue 
in  the  reduction  flask  in  which  the  iodine  was  added  before  the 


VANADIUM 


339 


bicarbonate,  the  over-reduction  averaging  0.0007  grm-  The 
determinations  by  reduction  in  the  residue,  in  which  the  neutral- 
ization took  place  before  the  addition  of  the  iodine,  uniformly 
show  an  incomplete  reduction  —  amounting  in  the  average  to 
0.0015  grm.  —  an  effect  which  is  without  doubt  due  to  the  action 
of  air  upon  the  sensitive  alkaline  solution  of  the  reduced  vanadate. 
Although  the  conditions  of  concentration  are  such  that  in  absence 
of  the  vanadic  acid  there  is  no  tendency  (barring  the  insignificant 
action  of  dissolved  air)  toward  liberating  iodine,  a  little  more 
iodine  is  liberated  by  vanadic  acid  when  acted  upon  by  sulphuric 
acid  and  potassium  iodide  than  would  correspond  to  a  reduction 
of  vanadic  acid  to  the  condition  of  the  tetroxide. 

Reduction  by  Hydrochloric  Acid  and  Potassium  Iodide  at  High  Concentrations. 


VsOg  in 
o.iooo  NH4V03 
taken. 

grm. 

HCl, 

concen- 
trated. 

cm.8 

KI. 

grm. 

Initial 
volume. 

cm.3 

Final 
volume. 

cm.3 

V203  found. 

Flask. 

Receiver. 

A. 


0.0632 

5 

0.6 

5° 

0.0323 

o  0632 

e 

i 

src 

2 

0.0613 

0.0627 
o  0327 

o  0632 

12  .  C, 

i 

CO 

2 

0.0618 

0.0637 
o  0378 

0.0632 
0.0632 

15 
25 

i 
i 

45 
50 

2 

* 
2 

* 
2 

0.0615 
0.0617 
0.0618 

o  .  0644 

0.0372 
0.0642 
0.0518 

0.0657 

Approximately  40  cm.3,  the  vanishing  point  of  the  iodine  color. 

B. 


0.0632 

15 

16 

2 

0.0618 

o  .  0630 

0.0632 

is 

16 

2 

0.0612 

0.0627 

0.0632 

15 

16 

2 

0.0617 

0.0625 

0.0632 

15 

16 

2 

0.0620 

o  .  0630 

0.0632 

15 

16 

2 

0.0616 

0.0627 

0.0632 

15 

16 

2 

0.0618 

0.0628 

0.0632 

15 

16 

2 

0.0617 

0.0630 

0.0632 

15 

16 

2 

0.0617 

o  0629 

0.0632 

15 

16 

2 

0.0616 

0.0629 

0.0632 

15 

16 

2 

0.0618 

o  .  0630 

On  the  other  hand,  the  degree  to  which  vanadic  acid  may  be 
reduced  by  hydrochloric  acid  and  potassium  iodide  is  found  to 
depend  upon  the  concentration.  By  carrying  the  distillation  to 


340 


METHODS  IN  CHEMICAL  ANALYSIS 


a  very  low  final  volume,  as  may  be  done  without  difficulty  in  the 
apparatus  described,  the  reduction  may  be  made  to  correspond 
very  closely  to  the  condition  of  vanadium  trioxide,  V2O3.  This 
is  shown  by  the  results  of  the  preceding  table. 

The  addition  of  phosphoric  acid,  as  proposed  by  Friedheim  and 
Euler  *  is  unnecessary  and  may  even  work  disadvantageously  by 
permitting  a  dangerous  rise  in  temperature  in  the  still  liquid 
residue  when  low  volumes  are  reached. 

This  is  shown  in  the  following  series  of  experiments  in  which 
I  grm.  of  potassium  iodide,  2  cm.3  of  sirupy  phosphoric  acid 
(sp.  gr.  1.70)  and  o.i  grm.  ammonium  vanadate  were  treated* 
trie  initial  volume  being  60  cm.3 

Reduction  by  Phosphoric  Acid  and  Potassium  Iodide  at  High  Concentrations. 


In  o.iooo  NH4VO3  taken. 

Final 
volume. 

cm.* 

Reduction  flask. 

Receiver. 

V204. 
grm. 

V203. 
grm. 

V204. 
grm. 

V203. 
grm. 

V204. 
grm. 

V203. 
grm. 

0.0699 
o  .  0699 
o  .  0699 
o  .  0699 
o  .  0699 
o  .  0699 
0.0699 
o  .  0699 
o  .  0699 

0.0632 
0.0632 
0.0632 
0.0632 
0.0632 
0.0632 
0.0632 
0.0632 
0.0632 

35 
25 

22 

4 
2 
2 

1-7 

1.6 
1.4 

0.0693 
0.0705 
0.0711 

0  .  0698 
0.0711 
o  .  0706 

o  .  0606 

* 

0.0597 

0.0621 
* 

o  .  0604 

0.0623 
0.0617 
0.0612 
o  0613 
0.0624 
0.0629 

*  Flask  broke. 

These  figures  indicate  that  when  the  distillation  is  continued 
until  the  volume  is  about  35  cm.3  the  condition  of  oxidation  cor- 
responds nearly  to  that  of  the  tetroxide  and  that  when  the  residue 
is  concentrated  almost  to  dryness  the  figures  approach  the  value 
for  the  trioxide.  Under  the  conditions  the  determinations  in 
the  residue  are  of  doubtful  value,  for  more  or  less  spattering 
occurs,  and  the  temperature  may  be  such  that  a  volatile  com- 
pound of  vanadium  begins  to  distil. 

It  is  evident,  therefore,  that  hydrochloric  acid  is  capable  of 

reducing  vanadic  acid,  even  in  the  cold,  to  the  condition  of  the 

tetroxide,  and  under  none  of  the  conditions  tried  does  reduction 

go  further;   that  hydrobromic  acid,  which  in  small  concentra- 

*  Ber.  Dtsch.  chem.  Ges.,  xxviii,  2071. 


VANADIUM  341 

tions  gives  a  definite  reduction  to  the  condition  of  the  tetroxide,* 
may  easily  push  the  reduction  well  on  toward  the  condition  of 
the  trioxide  at  higher  concentration;  and  that  the  reduction  by 
hydriodic  acid  may  be  carried  at  will  to  either  of  two  stages  — 
that  of  the  trioxide  or  that  of  the  tetroxide. 


The  Determination  of  Vanadic  Acid  by  Reduction  in  Acid  Solu- 
tion and  Reoxidation  by  Iodine  in  Alkaline  Solution. 

Reduction  by  A  method  for  the  determination  of  vanadium  has 
Organic  Acids.  Deen  described  by  Browning, f  in  which  tartaric 
acid  is  used  to  reduce  vanadic  acid  to  the  condition  of  the  tet- 
roxide which  is  then  reoxidized  in  alkaline  solution  by  standard 
iodine  and  thus  estimated.  Browning  and  Goodman  J  have 
shown  that  oxalic  acid  and  citric  acid  may  be  similarly  used  to 
reduce  vanadic  acid,  and  have  studied  the  conditions  under  which 
the  determination  is  accurate  in  presence  of  tungstic  and  molybdic 
acids. 

The  mode  of  proceeding  in  the  estimation  of  vanadium,  by 
the  use  of  either  tartaric,  oxalic  or  citric  acid  to  effect  reduction 
in  hot  solution,  may  be  briefly  summarized  as  follows:  To  the 
solution  of  a  vanadate  add  approximately  I  grm.  of  the  organic 
acid  for  every  tenth  of  a  gram  of  substance  to  be  determined. 
Heat  the  solution  to  boiling.  To  the  cooled  liquid  add  about 
5  grm.  of  potassium  bicarbonate  for  every  gram  of  acid  used. 
Add  iodine  in  slight  excess  and  set  aside  until  no  further 
bleaching  is  noticeable.  Destroy  the  excess  of  iodine  with 
standard  arsenite,  add  starch,  and  titrate  back  with  standard 
iodine.  The  total  amount  of  iodine  used  less  the  equivalent 
of  the  arsenious  oxide  is  the  measure  of  oxidation  according 
to  the  equation 

V2O4  +  I2  +  H2O  =  V205  +  2  HI. 

Results  of  this  procedure  are  given  on  pages  342,  343  and  344. 

The  process  of  reduction  by  means  of  tartaric  acid  in  boiling 

solution  fails  in  presence  of  molybdic  acid ,  which  under  the  con- 

*  Holverscheit's  procedure, 
t  Zeit.  anorg.  Chem.,  vii,  158. 

J  Philip  E.  Browning  and  Richard  J.  Goodman,  Am.  Jour.  Sci.,  [4],  ii, 
355- 


342 


METHODS  IN  CHEMICAL  ANALYSIS 


Reduction  by  Tartaric  Acid  in  Boiling  Solution:  Applicable  in  Presence  of 

Tungstic  Acid. 


V208  taken, 
grm. 

V2O5  found, 
grm. 

Error, 
grm. 

Sodium 
tungstate. 

grm. 

Tarraric  acid, 
grm. 

o.  1621 

0.1618 

—  0.0003 

2 

o.  1620 

o.  1624 

+0.0004 

2 

o.  1614 

0.1622 

+O.OOO8 

2 

o.  1619 

0.1606 

—  0.0013 

.  . 

I 

o.  1604 

0.1597 

—  0.0007 

.  . 

2 

0.1618 

0.1615 

—  0.0003 

3 

0.1298 

0.1305 

+0.0007 

i 

O.I2Q4 

0.1297 

+o  .  0003 

i 

0.1618 

0.1618 

0  .  0000 

2 

0.2588 

0.2575 

—0.0013 

3 

0.2722 

0.2726 

+0.0004 

.  . 

2 

0.3273 

0.3269 

—0.0004 

2 

0.1618 

0.1615 

—0.0003 

3 

o.  1615 

o.  1606 

—0.0009 

3 

0.1618 

0.1624 

+0.0006 

3 

0.1619 

o.  1624 

+0.0005 

3 

o.  1627 

0.1623 

—0.0004 

3 

0.1621 

o.  1624 

+0.0003 

4 

0.2587 

0.2574 

—0.0013 

4 

0.2587 

0.2589 

+O.OOO2 

4 

Reduction  by  Oxalic  Acid  in  Boiling  Solution:  Applicable  in  Presence  of  Tung- 
stic  Acid  and  Molybdic  Acid. 


V206  taken, 
grm. 

V2O5  found, 
grm. 

Error.  ' 
grm. 

Oxalic  acid, 
grm. 

Ammonium 
molybdate. 

grm. 

Sodium 

tungstate. 

grm. 

0.1806 

0.1803 

—  0.0003 

0.1950 

0.1955 

+o  .  0005 

.. 

o.i959 

0.1955 

—  0.0004 

o.  1950 

0.1959 

+0.0009 

o.i954 

0.1977 

+0.0023 

0.1956 

0.1960 

+O.OOO4 

.  . 

0.1956 

o.  1964 

+0.0008 

0.1956 

0.1957 

+O.OOOI 

0.3900 

0.3899 

—  O.OOOI 

2 

0.3897 

0.3917 

+O.O02O 

2 

0.3903 

o  .  3905 

+O.0002 

2 

0.1954 

0.1959 

+0.0005 

2 

I 

0.1957 

o.  1960 

+0.0003 

2 

I 

o.i954 

o.  1061 

+  O.OOO7 

2 

I 

0.1806 

0.1818 

+O.OOI2 

3 

o.  1807 

o.  1827 

+0.0020 

3 

0.1809 

0.1803 

—  0.0006 

3 

I 

o.  1956 

0.1961 

+o  .  0005 

3 

I 

0.3611 

0.3617 

+O.OOO6 

5 

0.3616 

0.3626 

+0.0010 

5 

I 

•  • 

VANADIUM 


343 


Reduction  by  Citric  Acid  in  Boiling  Solution:  Applicable  in  presence  of  Tung- 
stic  Acid  and  Molyjdic  Acid. 


V2OR  taken, 
grm. 

V2O6  found. 

gnu. 

Error, 
grm. 

Citric  acid, 
gnu. 

Ammonium 
molybdate. 

grm. 

Sodium 
tungstate. 

grm. 

0.1956 

0.1956 

O  .  0000 

I 

0-3905 

0.3921 

+0.0016 

2 

.     . 

0.1960 

o.  1960 

o  .  oooo 

I 

.    . 

0.1953 

o.  1060 

+0.0007 

I 

.  . 

. 

0.2088 

0.2082 

—  o  .  0006 

2 

_ 

O.  2100 

0.2098 

—  O.OOO2 

2 

.  . 

. 

o.  2092 

0.2107 

+0.0015 

I 

. 

o.  2092 

0.2107 

+0.0015 

2 

o.  2096 

0.2082 

—  0.0014 

2 

0-5 

0.2099 

o.  2116 

+0.0017 

3 

0-5 

o  .  2005 

0.2IOI 

+0.0006 

2 

o  5 

o  .  2099 

0.2095 

—  0.0004 

3 

.  .  . 

o  5 

ditions  is  also  somewhat  reduced.  Browning  and  Goodman  have 
shown,  however,  that  a  period  of  long  standing  —  twenty-four 
hours  —  may  be  substituted  in  this  process  for  the  period  of 
boiling  and  that  then  the  process,  otherwise  the  same,  is  effective 
in  presence  of  molybdic  acid  as  well  as  tungstic  acid. 
Reduction  by  Browning  *  has  shown  that  vanadic  acid  may  be 
Hydriodic  Acid.  recjucecl  by  the  action  of  potassium  iodidef  and  sul- 
phuric acid,  and  the  residue  of  vanadium  tetroxide  titrated  by 
iodine  in  alkaline  solution  {  with  sufficient  exactness  to  form  the 
basis  of  a  rapid  and  fairly  accurate  method  for  the  determination 
of  the  vanadium. 

According  to  this  procedure  a  vanadate  solution  is  placed  in 
a  trapped  Erlenmeyer  flask  §  with  potassium  iodide  (i  grm.  to 
2  grm.)  and  10  cm.3  of  sulphuric  acid  [i  :  i].  The  mixture  is 
boiled  until  fumes  of  iodine  are  no  longer  visible  and  the  escap- 
ing steam  gives  no  indication  of  free  iodine  with  red  litmus  paper.  || 
This  point  is  reached,  if  the  potassium  iodide  is  kept  within  the 
above  limits,  when  the  volume  has  diminished  from  approxi- 
mately 100  cm.3  to  35  cm.3  The  cooled  solution  is  nearly  neu- 

*  Philip  E.  Browning,  Am.  Jour.  Sci.,  [4],  ii,  185. 

f  See  page  337. 

J  Browning,  Zeit.  anorg.  Chem.,  vii,  158. 

§  See  Fig.  6,  page  6. 

II  See  page  450. 


344 


METHODS  IN  CHEMICAL  ANALYSIS 


Reduction  by  Tar  tar  ic  Acid  on  Standing  24  Hours  or  More:  Applicable  in 
Presence  of  Tungstic  Acid  and  Molybdic  Acid. 


VA  taken, 
grm. 

V208  found, 
grm. 

Error, 
grm. 

Time  in. 
days. 

Tartaric  acid, 
grm. 

Total 
volume. 

cm.1 

0.1646 

O.  1649 

+o  .  0003 

I 

4 

25 

0.1640 

O.  1606 

-0.0034 

I 

4 

65 

0.1293 

O.  1264 

—  0.0029 

2 

3 

55 

0.1633 

0.1628 

-0.0005 

2 

4 

65 

0.1293 

0.1288 

—  0.0005 

3 

2-5 

50 

o.  1298 

0.1299 

+O.OOOI 

3 

2-5 

50 

0.1295 

0.1279 

—  0.0016 

3 

3 

55 

0.1617 

0-1597 

—  O.OO2O 

4 

2 

70 

0.1623 

0.1622 

—  o.oooi 

4 

3 

80 

taken, 
grm. 

VA 
found. 

grm. 

Error, 
grm. 

Ammonium 
molybdate. 

grm. 

Sodium 
tungstate. 

grm. 

Time  in 
days. 

Total 
volume. 

cm.* 

Tartaric 
acid. 

grm» 

0.1552 

0.1558 

+O  .  0006 

I 

25 

5 

0.1289 

0.1301 

+  O.OOI2 

. 

.  . 

I 

25 

5 

0.2583 

0.2587 

+0.0004 

.  .  . 

50 

5 

0.1293 

0.1299 

0.0006 

I 

25 

6 

0.2582 

0.2591 

+0.0009 

I 

50 

6 

0.2582 

0.2588 

+O.OOO6 

I 

50 

5 

0.1297 

0.1308 

+O.OOII 

I 

.  . 

25 

5 

o.  1291 

0.1289 

—  0.0002 

.  .  . 

I 

25 

6 

0.2582 

0.2568 

—0.0014 

I 

50 

5 

0.1293 

0.1299 

+0.0006 

I 

I 

I 

25 

8 

0.2582 

0.2579 

—0.0003 

I 

I 

I 

50 

5 

0.1550 

0.1538 

—  O.OOI2 

2 

25 

5 

0.1556 

0.1545 

—  O.OOII 

2 

25 

5 

0.1289 

0.1296 

+0.0007 

2 

25 

5 

0-1549 

0.1527 

—  O.O022 

0-5 

2 

25 

5 

0.1553 

0.1548 

—  0.0005 

I 

2 

25 

5 

0.1556 

0.1554 

—  O.OOO2 

I 

2 

25 

5 

0.1293 

0.1310 

+0.0017 

I 

2 

25 

6 

0.1295 

0.1299 

+0.0004 

.  .  . 

I 

2 

25 

6 

0.1293 

0.1289 

—  O.OOO4 

I 

I 

2 

25 

7 

0.1293 

0.1301 

+0.0008 

3 

25 

5 

o.  1289 

0.1299 

+O.OOIO 

0-5 

3 

25 

5 

0.1293 

o.  1292 

—  O.OOOI 

I 

3 

25 

7 

0.1556 

0.1567 

+O.OOII 

I 

3 

30 

5 

o.  1291 

o.  1289 

—  O.OOO2 

I 

I 

3 

25 

7 

0.1550 

0.1557 

+0.0007 

4 

25 

5 

0.1554 

0.1557 

+0.0003 

I 

4 

25 

5 

0.1556 

0.1557 

+0.0001 

0-5 

4 

25 

5 

VANADIUM 


345 


tralized  by  sodium  hydroxide,*  a  little  tartaric  acid  is  added  to 
prevent  precipitation  of  the  tetroxide,  and  neutralization  is  com- 
pleted by  acid  potassium  carbonate.  To  the  cooled  solution  w/io 
iodine  is  added  in  slight  excess,  the  flask  is  closed  with  a  paraffined 
cork,  and  the  mixture  is  allowed  to  stand  about  half  an  hour.  The 
excess  of  iodine  is  taken  up  by  n/io  arsenite,  and  the  excess  of 
the  last  is  titrated  with  n/io  iodine  in  presence  of  starch.  The 
difference  between  the  amount  of  n/io  arsenite  used,  expressed  in 
terms  of  iodine,  and  the  whole  amount  of  n/io  iodine  employed 
gives  the  amount  of  iodine  necessary  to  oxidize  the  vanadium 
tetroxide  to  vanadium  pentoxide. 

Results  of  this  procedure  are  given  in  the  table. 

Reduction  by  Sulphuric  Acid  and  Potassium  Iodide:  Reoxidation  by  Iodine  in 
Alkaline  Solution:    Final  Titration  with  Standard  Arsenite. 


V2O5  taken, 
grm. 

V2O6  found, 
grm. 

Error, 
grm. 

KI. 

grm. 

H2SO4fi  :  ij. 
cm.3 

0.1699 

o.  1690 

—  0.0009 

IO 

o  .  i  704 

0.1699 

—  0.0005 

IO 

o  .  i  706 

o.  1700 

—  0.0006 

10 

0.1702 

0.1692 

—  0.0010 

TO 

0.3613 

o  3620 

+0.0007 

•  5 

10 

o.  1805 

o  1803 

+O.OOO2 

10 

0.3614 

o  3620 

-|-O.OOO6 

•  5 

IO 

o.  1811 

o.  1814 

+0.0003 

IO 

o.  1807 

o  1815 

+0.0008 

IO 

0.3613 

0.3620 

+O.OO07 

•5 

IO 

0.3679 

0.3674 

—  0.0005 

•  5 

IO 

0.3612 

0.3608 

—  O.OOO4 

•  5 

IO 

0.2893 

0.2907 

—  O.OOI4 

•  5 

IO 

0.3456 

0.3448 

—  0.0008 

•  5 

10 

0-3453 

0.3448 

—  O.OOO5 

•  5 

10 

0.3907 

0.3912 

+O.OOO5 

2 

10 

0.3908 

0.3898 

—  o.ooio 

I 

IO 

0.3906 

0.3921 

+0.0015 

2 

IO 

0.3909 

0.3912 

+o  .  0003 

i-5 

IO 

Reduction  by 
Hydrobromic 
Acid. 


For  the  estimation  of  the  bromine  set  free  by  the 
action  of  potassium  bromide  and  sulphuric  acid  upon 
a  vanadate  according  to  Holverscheit,  Browning  f 
has  substituted  the  reoxidation  of  the  reduced  vanadium  tet- 
roxide by  iodine  in  presence  of  an  acid  carbonate. 

*  The  potassium  or  sodium  hydroxide  for  this  work  must  be  free  from  alco- 
hol, a&  the  solution  is  allowed  to  stand  with  iodine  after  neutralization.  It 
is  conveniently  prepared  by  mixing  potassium  or  sodium  carbonate  in  proper 
proportions  with  calcium  oxide  and  filtering  off  the  calcium  carbonate. 

t  Philip  E.  Browning,  Am.  Jour.  Sci.,  [4],  ii,  185. 


346 


METHODS  IN  CHEMICAL  ANALYSIS 


According  to  this  method,  reduction  is  effected  by  boiling  in 
a  trapped  Erlenmeyer  flask  *  a  mixture  of  the  vanadate  in  solu- 
tion with  I  grm.  to  2.5  grm.  of  potassium  bromide  and  10  cm.3 
of  [i  :  i]  sulphuric  acid  until  bromine  is  no  longer  present  in  the 
steam,  a  point  ordinarily  reached  when  the  original  volume  of 
about  100  cm.3  has  diminished  to  a  volume  of  about  25  cm.3. 
At  this  point  the  solution  is  nearly  neutralized  with  alkali 
hydroxide  free  from  alcohol  f  and  completely  neutralized  by  an 
acid  alkali  carbonate,  a  little  tartaric  acid  having  been  first  added 
to  prevent  precipitation  of  the  vanadium  tetroxide.  An  excess  of 
n/io  iodine  is  added  and  the  mixture  is  allowed  to  stand  half  an 
hour  with  the  flask  closed  by  a  paraffined  stopper.  The  excess  of 
iodine  is  taken  up  by  n/io  arsenite,  and  the  slight  excess  of  the 
last  is  determined  by  titration  with  n/io  iodine.  The  difference 
between  the  entire  amount  of  iodine  used  and  the  iodine  equiva- 
lent of  the  arsenite  employed  measures  the  oxidation  of  vanadium 
tetroxide  to  vanadium  pentoxide. 

Results  of  this  procedure  are  given  in  the  table. 

Reduction  by  Sulphuric  Acid  and  Potassium  Bromide:    Reoxidation  by  Iodine 
in  Alkaline  Solution:   Titration  with  Arsenite. 


V2O5  taken, 
grm. 

V2O6  found, 
grm. 

Error, 
grm. 

KBr. 

grm. 

H2S04[i:i|. 
cm.8 

o.  1800 

0.1876 

—  O.OOI4 

I 

10 

0.1886 

0.1886 

O.OOOO 

2 

IO 

0.1885 

0.1882 

—  0.0003 

I 

IO 

0.1885 

0.1886 

+O.OOOI 

i-S 

IO 

0.1881 

0.1873 

—0.0008 

i-5 

10 

0.1886 

0.1882 

—0.0004 

2 

IO 

0.3907 

0.3804 

—0.0013 

2 

IO 

0.3907 

0.3903 

—0.0004 

2 

IO 

0.3907 

0.3894 

—0.0013 

2 

10 

0.3909 

0.3889 

—  O.OO2O 

2 

IO 

0.3911 

0.3903 

—  0.0008 

1-5 

IO 

0.3902 

0.3900 

—  O.OOO2 

2-5 

IO 

The  Use  of  the  Jones  Reductor  in  the  Estimation  of  Vanadic  Acid. 

The  reduction  of  vanadic  acid  to  the  condition  of  vanadium 

tetroxide  preparatory  to  estimating  it  by  titration  with  potassium 

permanganate  may  be  accomplished  by  the  use  of  sulphurous 

acid  or  hydrogen   sulphide.     The  more  convenient  method  of 

*  See  Fig.  6,  page  6.  t  See  note,  page  345. 


VANADIUM  347 

treatment  by  zinc  and  free  acid,  followed  by  direct  titration,  is  not 
applicable  on  account  of  the  irregularities  in  reduction  found 
under  the  conditions.  Gooch  and  Gilbert  *  have  developed  a 
reliable  method  for  bringing  the  product  of  reduction  of  vanadic 
acid  by  zinc  and  acid  definitely  to  the  condition  of  the  tetroxide, 
in  order  that  the  Jones  reductor  so  useful  in  the  preparation  of 
salts  of  iron  for  titration  by  potassium  permanganate  may  be 
made  similarly  applicable  in  the  estimation  of  vanadic  acid  and 
its  salts. 

A  convenient  form  of  the  Jones  reductor  f  is  made  as  follows : 
The  contracted  end  of  a  piece  of  glass  tubing,  2  cm.  in  inside 
diameter  and  50  cm.  long,  is  sealed  to  a  stopcock  prolonged  in 
a  smaller  tube,  0.5  cm.  in  inside  diameter,  to  a  length  of  24  cm. 
At  the  point  of  contraction  in  the  larger  tube  is  placed  a  piece 
of  platinum  gauze,  next  to  this  a  mat  of  fine  glass  wool  2  cm. 
in  thickness,  and  upon  the  last  a  column  40  cm.  long  of  amal- 
gamated zinc  of  a  size  to  pass  a  sieve  of  eight  meshes  to  the 
centimeter.  The  smaller  tube  is  passed  through  a  rubber  stopper 
fitted  to  a  vacuum  flask  and  the  last  is  connected  through  a 
pressure  regulator  with  the  vacuum  pump.  In  using  this  appa- 
ratus the  pump  is  started,  the  regulator  set  to  give  a  pressure 
in  the  flask  less  than  the  outside  pressure  by  an  amount  equal 
to  20  cm.  of  water,  the  reducing  column  of  zinc  warmed  by  pass- 
ing through  it  hot  distilled  water  followed  by  100  cm.3  of  hot 
I  per  cent  sulphuric  acid,  and  then  the  solution  of  the  salt  of 
vanadic  acid  to  be  reduced  is  gradually  drawn  through  the  zinc 
in  small  portions,  alternating  with  portions  of  the  I  per  cent 
sulphuric  acid  amounting  to  100  cm.3  Finally,  the  column  is 
washed  down  with  100  cm.3  more  of  the  dilute  sulphuric  acid  and 
about  250  cm.3  of  hot  distilled  water.  Throughout  the  entire 
series  of  operations  care  is  taken  to  keep  the  zinc  covered  with 
liquid  and  so  out  of  contact  with  the  air. 

The  lavender  solution  collected  in  the  flask  contains  vanadium 
dioxide,  which,  when  exposed  to  the  action  of  a  current  of  air 
takes,  as  Roscoe  has  shown, {  the  blue  color  of  the  tetroxide; 
but  the  action  of  the  molecular  oxygen  of  the  air  proves  to  be 
insufficient  to  bring  about  the  complete  oxidation  of  the  lower 

*  F.  A.  Gooch  and  R.  D.  Gilbert,  Am.  Jour.  Sci.,  [4],  xv,  389. 
t  Described  in  Blair's  Chemical  Analysis  of  Iron, 
t  Ann.  Pharm.  Suppl.,  vi,  98  (1868). 


348 


METHODS  IN  CHEMICAL  ANALYSIS 


vanadium  oxides  in  acid  solution  to  the  condition  of  the  tetroxide 
within  a  reasonable  time.  Ordinary  oxidizers  carry  the  action 
too  far,  oxidizing  the  tetroxide  as  well  as  the  lower  oxides.  It  is 
found,  however,  that  silver  oxide  and  silver  salts  supply  oxygen 
sufficiently  to  affect  the  lower  oxides  easily  while  leaving  the 
tetroxide  intact.  Silver  sulphate  appears  to  be  the  most  con- 
venient form  in  which  to  use  the  silver  compound. 
Regulation  of  In  treating  the  solution,  therefore,  in  the  reductor 
Use  of  suve?  as  described,  the  receiving  flask  is  charged  with 
Sulphate.  a  saturated  solution  of  silver  sulphate,  and  when  the 

reduced  solution  issuing  from  the  reductor  meets  the  silver  sul- 
phate a  muddy  deposition  of  finely  divided  silver  begins.  The 
mixture  is  boiled  and  filtered  upon  asbestos  in  a  perforated 
crucible.  The  solution,  now  about  700  cm.3  in  volume,  is  heated 
to  the  boiling  point  and  titrated  with  n/io  potassium  perman- 
ganate. Upon  boiling  the  mixture,  the  metallic  silver  gathers 
into  a  single  spongy  mass  and  leaves  the  solution  so  clear  that, 
were  it  not  that  spongy  silver  is  easily  acted  upon  by  the  per- 
manganate, the  titration  might  be  made  without  previous 
nitration. 

For  small  amounts  of  vanadium  the  determinations  are 
accordant  and  exact.  Wider  variations,  due  to  the  difficulty  of 
catching  the  pink  end  reaction  in  presence  of  the  reddish  yellow 
color  which  appears  as  the  vanadic  acid  is  formed  in  consider- 
able amount,  are  inherent  in  the  permanganate  process  of  titra- 
tion when  large  amounts  of  vanadic  acid  are  involved. 

The  results  of  test  experiments  follow  in  the  table : 


Reduction  by  Zinc:    Partial   Oxidation  by  Silver  Sulphate: 

Permanganate. 


Titration  -with 


V2O5  taken. 

KMnO4  required 
nearly  n/2o. 

V2O5  found. 

Error. 

grm. 

cm.8 

grm. 

grm. 

0.0767 

17 

0.0770 

+0.0003 

0.0767 

17.04 

0.0771 

+O  .  0004 

0.0767 

I7-05 

0.0772 

+0.0005 

0.0767 

16.96 

0.0768   . 

+O.OOOI 

0.0767 

16.98 

0.0769 

+O.OO02 

0.0767 

17 

0.0770 

+0.0003 

0.1918 

42-9 

0.1942 

+0.0024 

0.1918 

42.7 

0-1933 

+0.0015 

VANADIUM 


349 


Registration  of  In  a  study  of  the  effects  of  zinc  and  magnesium  in 
ust^Ferric  tne  reduction  of  vanadic  acid,  Gooch  and  Edgar,* 
Sulphate.  have  shown  that  the  degree  in  which  vanadic  acid 

may  be  reduced  by  magnesium  in  the  presence  of  hydrochloric 
acid  or  sulphuric  acid  is  irregular  and  dependent  upon  conditions 
not  easily  controlled.  With  magnesium  amalgam  the  reduction 
proceeds  most  readily  and  approximates  more  or  less  to  the  condi- 
tion of  V2O2,  but  under  none  of  the  conditions  tried  is  the  reduc- 
tion by  magnesium  sufficiently  definite  to  be  applied  in  a  good 
analytical  process.  With  zinc  the  reduction  is  more  regular, 
but  the  condition  of  oxidation  of  the  product  of  the  reduction  of 
vanadic  acid  by  zinc  and  hydrochloric  acid  in  the  flask,  by  zinc 
and  sulphuric  acid  in  the  flask,  or  by  amalgamated  zinc  in  the 
reductor,  never  corresponds  exactly  to  V2O2  when  titration  is 
made  in  air.  It  is  shown  that  the  use  of  the  zinc  reductor 
carries  the  reduction  easily  and  rapidly  to  the  condition  of  V2O2, 
and  that  by  anticipating  the  oxidizing  action  of  the  air  by  means 
of  a  ferric  salt  in  the  receiver  f  the  solution  is  made  less  sensitive 
to  the  action  of  the  air,  while  the  highest  degree  of  reduction  is 
registered  by  the  ferrous  salt  formed;  but  a  solution  containing 
vanadium  in  the  condition  of  V2O2  cannot  be  exposed  to  air, 
even  momentarily,  without  undergoing  oxidation. 

Reduction  by  Zinc:    Treatment  with  Ferric  Salt:    Titration  by  Permanganate. 


Ferric 
alum, 
10  per  cent 
solution. 

Phosphoric 
acid  syrup. 

KMnO4 
«/ioX  1.052. 

V2O5  taken 
as  NaV03. 

V2O5  calculated 
from  oxidation 
from  V2O2. 

Error  in 
terms  of 
V205. 

cm.* 

cm.3 

cm.8 

grm. 

grm. 

grm. 

25 

5 

43-10 

0.1381 

0.1378 

—  0.0003 

25 

5 

43.20 

0.1381 

0.1381 

o.oooo 

25 

5 

43-30 

0.1381 

0.1384 

+0.0003 

25 

5 

43-28 

0.1381 

o.  1384 

+0.0003. 

25 

5 

43-32 

0.1381 

0.1385 

+0.0004 

15 

3 

21.  60 

0.0691 

o  .  0690 

—  o.ooor 

15 

3 

21  .62 

0.0691 

0.0691 

o.oooo 

IS 

3 

21.  80 

0.0691 

o  .  0696 

+0.0005 

40 

8 

64.90 

0.2072 

0.2075 

+0.0003 

According  to  the  recommended  procedure,  the  receiver  at- 
tached to  the  zinc  column  is  charged  with  a  solution  of  ferric 

*  F.  A.  Gooch  and  Graham  Edgar,  Am.  Jour.  Sci.,  [4],  xxv,  233. 
t  See  page  426. 


35°  METHODS  IN  CHEMICAL  ANALYSIS 

alum,  gentle  suction  is  applied,  and  through  the  column  of  amal- 
gamated zinc  are  passed  in  succession  hot  water  (100  cm.3),  2.5  per 
cent  sulphuric  acid  (100  cm.3),  the  solution  of  vanadic  acid  in 
2.5  per  cent  sulphuric  acid  (125  cm.3),  and  finally  hot  water 
(200  cm.3).  To  the  receiver  is  added  sirupy  phosphoric  acid 
(4  cm.3),  to  decolorize  the  solution,  and  the  titration  of  the  hot 
solution  is  made  in  the  usual  manner  with  potassium  perman- 
'ganate.  Correction  should  be  made  for  the  action  of  the  zinc 
column  upon  the  reagents  without  the  vanadic  acid.  Experi- 
mental tests  of  the  method  are  given  in  the  table. 

The  Estimation  of  Vanadic  and  Arsenic  Acids  and  of  Vanadic  and 
Antimonic  Acids  in  the  Presence  of  One  Another. 

If  a  solution  containing  arsenic  and  vanadic  acids  or  antimonic 
and  vanadic  acids  is  boiled  with  tartaric  or  oxalic  acids,  the 
vanadic  acid  is  reduced  to  tetroxide  and  may  be  reoxidized  in 
alkaline  solution  by  iodine  *  according  to  the  equation 

V2O4  +  I2  +  H2O  =  V2O5  +  2  HI. 

If  a  solution  containing  arsenic  and  vanadic  acids,  or  antimonic 
and  vanadic  acids,  is  reduced  with  sulphur  dioxide  under  proper 
conditions,  the  arsenic  acid  is  reduced  to  arsenious  acid,  or  the 
antimonic  acid  to  antimonious  acid,  and  the  vanadic  acid  to  tet- 
roxide, so  that  after  boiling  off  the  excess  of  reagent  the  reoxi- 
dation  by  iodine  in  alkaline  solution  should  proceed  according 
to  one  of  the  eolations 

As203  +  V204  +  3  Ia  +  3  H20  =  As205  +  V2O5  +  6  HI 
Sb203  +  V204  +  3  I2  +  3  H20  =  Sb205  +  V205  +  6  HI, 

and  the  iodine  thus  used  should  in  each  case  correspond  to  the 
sum  of  the  two  oxides.  Edgar  f  has  shown  that  if  aliquot  parts 
of  the  same  solution  are  treated  according  to  these  processes 
the  titration  of  the  solution  reduced  by  tartaric  acid  determines 
the  vanadium  present  and  that  the  titration  of  the  solution 
reduced  by  sulphur  dioxide  determines  the  sum  of  the  oxides  of 
arsenic  and  vanadium,  or  of  antimony  and  vanadium,  the  differ- 
ence in  the  number  of  cubic  centimeters  of  iodine  used  in  the  two 
*  Browning,  Zeit.  anorg.  Chem.,  vii,  158;  Browning  and  Goodman,  Am. 
Jour.  Sci.,  [4],  ii,  355- 

t  Graham  Edgar,  Am.  Jour.  Sci.,  [4],  xxvii,  299. 


VANADIUM 


351 


cases  being  the  amount  required  to  oxidize   the  arsenic  oxide 
or  antimony  oxide. 

The  method  of  treatment  in  detail  is  as  follows:  The  solu- 
tion of  arsenate  and  vanadate,  or  of  antimonate  and  vanadate, 
is  divided  into  two  portions. 

(I)  One  portion  is  boiled  with  I  to  2  grm.  of  tartaric  or  oxalic 
acid  until  the  blue  color  of  the  vanadium  tetroxide  indicates 
complete  reduction.     The  solution  is  then  cooled,  nearly  neu- 
tralized with  potassium  bicarbonate,  and  an  excess  of  standard 
iodine  solution  is  added.     Neutralization  is  completed,  an  excess 
of  bicarbonate  added,  and  the  solution  allowed  to  stand  for  from 
fifteen  to  thirty  minutes.     The  excess  of  iodine  is  then  removed 
with  standard  arsenious  acid  and  the  solution  titrated  to  color 
after  the  addition  of  starch. 

(II)  The  second  portion  of  the  solution  is  placed  in  a  small 
pressure   flask   and   slightly   acidified   with   sulphuric   acid.     A 
strong  solution  of  sulphurous  acid  (25  cm.3)  is  then  added  and 
the  flask  is  closed  and  heated  for  one  hour  on  the  steam  bath.* 
After  cooling,  the  flask  is  opened  and  the  solution  transferred  to 
an  Erlenmeyer  flask  and  boiled  to  remove  the  excess  of  sulphur 
dioxide,  a  current  of  carbon  dioxide  being  passed  into  the  liquid 
to  facilitate  the  removal  of  the  last  traces.     The  solution  is  then 
cooled,  nearly  neutralized  with  potassium  bicarbonate  and  an 


Arsenic  and  Vanadium. 


VA 

taken. 

V205 
found. 

V206 
error. 

As205 
taken. 

As205 
found. 

•     As205 
error. 

(I) 

n/io 
iodine. 

(ID 

n/io 
iodine. 

giro. 

grm. 

grm. 

grm. 

grm. 

grm. 

cm.* 

cm.1 

0.1183 

0.1181 

—  O.0002 

o  .  0960 

0.0961 

+O.OOOI 

12-95 

29.65 

0.1183 

o.  1183 

0.0000 

O  .  0960 

0.0962 

+O.OOO2 

12.97 

29.70 

0.1183 

0.1182 

—  O.OOOI 

O  .  0960 

0.0962 

+O.OOO2 

12.96 

29.70 

0.0591 

0.0593 

+O  .  OOO2 

O  .  0480 

o  .  0480 

O.OOOO 

6.50 

14-85 

9-0591 

0.0594 

+0.0003 

O  .  0480 

0.0482 

+O.OOO2 

6.52 

14.90 

0.0591 

0.0589 

—  O.OOO2 

O  .  0480 

o  .  0483 

+0.0003 

6.45 

14.86 

0.1774 

0.1779 

+o  .  0005 

0.1440 

0.1438 

—  O.OOO2 

19.50 

44-50 

0.1774 

0.1774 

O  .  OOOO 

O.  1440 

0.1440 

O.OOOO 

19-45 

44-50 

0.1774 

0.1776 

+0.0002 

O.  1440 

0.1442 

+O.OOO2 

19-47 

44-45 

0.2366 

0.2371 

+0.0005 

o  .  0480 

o  .  0478 

—  O.OOO2 

26.00 

34-31 

0.2366 

0.2366 

o.oooo 

o  .  0480 

o  .  0480 

O.OOOO 

25-95 

34.30 

0.0591 

0-0593 

+O  .  OOO2 

o.  1440 

0.1439 

—  O.OOOI 

6.50 

38-90 

*  McCay,  Am.  Chem.  Jour.,  vii,  273;  Von  Knorre,  Zeit.  angew.   Chem. 
(1888),  155. 


352 


METHODS   IN   CHEMICAL   ANALYSIS 

Antimony  and  Vanadium. 


-     (I) 

(EH 

V205 

V205 

V206 

Sbo05 

Sb205 

Sb20e 

«/ioX 

w/ioX 

taken. 

found. 

error. 

taken. 

found. 

error. 

0.9375 

p.9375 

iodine. 

iodine. 

gnu. 

grm. 

grm. 

grm. 

grm. 

grm. 

cm8. 

cms. 

0.1183 

0.1185 

+O.OO02 

0.0757 

0.0759 

+  O.OOO2 

13-85 

23.69 

0.1183 

0.1186 

4-o.ooo3 

0.0757 

0.0764 

+O.OOO7 

13-87 

24.05 

0.1183 

0.1183 

o.oooo 

0.0757 

0.0760 

+0.0003 

13-83 

23-95 

0.1774 

0.1777 

+0.0003 

0.1261 

0.1258 

—  0.0003 

20.80 

37-55 

0.1774 

0.1777 

+0.0003 

0.1261 

0.1258 

—  0.0003 

20.80 

37-55 

0.1774 

0.1773 

—  o.oooi 

0.1261 

0.1260 

—o.oooi 

20.75 

37-50 

0.2366 

0.2376 

+O.OOIO 

0.1261 

0.1257 

—  0.0004 

27.80 

44-52 

0.2366 

0.2369 

+0.0003 

0.1261 

0.1263 

—  O.OOO2 

27.70 

44-50 

0.2366 

0.2369 

+0.0003 

0.1261 

0.1260 

—o.oooi 

27.70 

44-45 

excess  of  iodine  added  as  before.  After  completing  the  neutral- 
ization, adding  an  excess  of  bicarbonate  and  allowing  to  stand 
for  one-half  hour,  the  excess  of  iodine  was  determined  by  arse- 
nious  acid  as  before. 

As  before  stated,  the  titration  figures  of  (I)  determine  the 
vanadium,  and  the  subtraction  of  these  figures  from  those  of 
(II)  gives  figures  for  the  arsenic  or  the  antimony,  as  the  case 
may  be. 

Results  of  the  method  are  given  in  the  table. 

The  Estimation  of    Vanadic  Acid   Associated  with   Chromium, 
with  Molybdenum,  and  with  Iron. 

Methods  for  the  estimation  of  vanadic  acid  associated  with 
chromic  acid,  with  a  ferric  salt,  with  chromic  acid  and  a  ferric 
salt,  or  with  molybdic  acid,  are  described  elsewhere.* 

The  Estimation  of  Vanadium  in  the  Tetroxide  Condition  by  the 

Action  of  Potassium  Ferricyanide  in  Alkaline  Solution 

and  Potassium  Permanganate  in  Acid  Solution. 

To  the  estimation  of  vanadium  in  the  tetroxide  condition  of 
oxidation,  Palmer  f  has  applied  the  oxidizing  action  of  potassium 
ferricyanide  in  alkaline  solution  with  reoxidation  of  the  resulting 
ferrocyanide  in  acid  solution,  following  the  general  plan  formerly 
employed  by  Browning  and  Palmer  for  the  estimation  of  cerium  t 

*  See  page  510. 

t  Howard  E.  Palmer,  Am.  Jour.  Sci.,  [4],  xxx,  141. 

t  See  page  249. 


VANADIUM 


353 


and  thallium,*  and  by  Palmer  for  the  estimation  of  arsenic, 
antimony  and  tint-  The  reactions  involved  may  be  repre- 
sented by  the  following  equations: 

V2O4  +  2  K3FcC6N6  +  2  KOH  =  V2O5  +  2  K4FeC6N6  +  H2O, 
5  K4FeC6N6  +  KMnO4  +  4  H2SO4  = 

5  KsFeCeNe  +  3  K2SO4  +  MnSO4  +  4  H2O. 

Vanadium  was  prepared  in  the  tetroxide  condition  by  treating 
portions  of  a  solution  of  ammonium  vanadate,  made  slightly 
acid  with  hydrochloric  acid,  with  a  current  of  sulphur  dioxide 
until  the  clear  blue  color  indicated  complete  reduction  of  the 
vanadic  acid  to  the  condition  of  the  tetroxide,  V2O4.  The  solu- 
tion was  then  boiled  in  a  current  of  carbon  dioxide  until  the  last 
traces  of  sulphur  dioxide  had  been  removed.  To  determine  the 
vanadium  in  such  a  solution  the  following  procedure  was  adopted. 

To  the  cooled  solution  are  added  in  solution  potassium  ferri- 
cyanide  to  at  least  tenfold  the  amount  theoretically  necessary 
for  the  oxidation  and  about  6  grm.  of  potassium  hydroxide, 
the  concentrations  being  such  that  the  total  volume  amounts 
to  100  cm.3  or  125  cm.3.  Barium  hydroxide  is  added  to  precipi- 
tate the  vanadate,  which  otherwise  would  form  with  the  ferro- 
cyanide  J  an  insoluble  compound  refractory  in  the  subsequent 
titration  by  the  permanganate.  The  precipitate  is  settled  and 
filtered  off  on  an  asbestos  felt.  The  filtrate  and  washings  are 
acidified  with  hydrochloric  acid  and  titrated  with  permanganate. 
The  results  obtained  are  recorded  in  the  table. 

Estimation  of  Vanadium  Tetroxide. 


V2O6  taken.* 
grm. 

K3PeC6N6  used, 
grm. 

KOH  used, 
grm. 

V?O5  found, 
grm. 

Error, 
grm. 

o  .  0060 

4 

6 

0.0959 

—  o.oooi 

o  .  0960 

4 

6 

0.0954 

—0.0006 

O  .  OQ^O 

4 

6 

0.0956 

—  0.0004 

o  .  0960 

4 

6 

0.0962 

+  O.OOO2 

o  .  0960 

4 

6 

0.0956 

—  o  .  0004 

o  .  0960 

4 

6 

o  .  0959 

—  O.OOOI 

o  .  0960 

4 

6 

0.0961 

-f-o  .  oooi 

o  .  0960 

4 

6 

0.0962 

-j-o.oooi 

o  .  0960 

4 

6 

0.0961 

o.oooo 

0.0960 

4 

6 

0.0961 

+0.0001 

*  Reduced  by  sulphur  dioxide  as  described. 

*  See  page  223.  f  See  page  322. 

J  Griitzner,  Chem.  Centralblatt,  1902,  i,  500. 


354  METHODS  IN  CHEMICAL  ANALYSIS 

In  this  process  the  large  proportions  of  ferricyanide  and 
potassium  hydroxide  are  necessary  to  insure  complete  oxidation 
of  the  vanadium  at  the  concentrations  employed.  If  the  dilu- 
tion is  greater,  more  ferricyanide  is  required.  Titration  by 
permanganate  in  a  solution  acidified  with  sulphuric  acid  is 
unsatisfactory  on  account  of  the  difficulty  in  noting  the  end  point 
in  the  presence  of  barium  sulphate  precipitated  by  the  action  of 
the  excess  of  barium  hydroxide.  The  ferrocyanide,  however, 
may  be  safely  titrated  by  permanganate  in  the  cold  in  dilute 
hydrochloric  acid  solution. 


CHAPTER   IX. 
OXYGEN;  SULPHUR;  SELENIUM;  TELLURIUM. 

OXYGEN. 

The  lodometric  Determination  of  Oxygen  in  Air  and  in  Aqueous 

Solution. 

THE  reactions  by  which  a  satisfactory  iodometric  determina- 
tion of  the  oxygen  of  perchlorates  may  be  accomplished  *  have 
also  been  applied  by  Kreider  f  to  the  determination  of  the 
oxygen  of  the  air  or  of  oxygen  dissolved  in  water. 
Determination  of  The  method,  in  brief,  consists  in  conducting  a 
oxygen  in  Air.  known  volume  of  air  through  a  strong  solution  of 
hydriodic  acid  in  the  presence  of  nitrogen  dioxide,  subsequently 
neutralizing  the  acid  with  potassium  bicarbonate,  titrating  the 
liberated  iodine  with  standard  decinormal  arsenic  solution,  and 
calculating  the  equivalent  weight  and  then  the  volume  of  oxygen. 
By  several  simple  devices,  to  be  described,  all  calculations  may 
be  done  away  with  and  the  percentage  of  oxygen  seen  immedi- 
ately by  the  volume  of  arsenic  solution  required  for  titration. 

The  volume  of  oxygen  indicated  under  the  standard  condi- 
tions of  temperature  and  pressure  (o°  and  760  mm.)  must  either 
be  calculated  to  that  which  it  would  occupy  under  the  conditions 
of  the  experiment,  or  the  volume  of  air  taken  must  be  reduced 
to  the  standard  conditions  of  temperature  and  pressure.  The 
latter  plan  is  the  more  satisfactory,  since  by  Lunge's  ingenious 
device  {  the  reduction  can  be  readily  effected  without  any 
calculation,  and  independently  of  changing  temperature  and  pres- 
sure. For  this  purpose  the  following  arrangement  of  two  bu- 
rettes answered  admirably.  One  burette  graduated  to  120  cm.3 
contains  over  mercury  the  same  volume  of  moist  air  which 
100  cm.3  of  moist  air  at  o°  and  760  mm.  would  occupy  under  the 
given  conditions,  this  standard  being  very  carefully  de- 

*  See  page  467. 

t  D.  Albert  Kreider,  Am.  Jour.  Sci.,  [4],  ii,  361. 
t  Zeit.  angew.  Chem.,  1890,  139. 
355 


356 


METHODS  IN  CHEMICAL  ANALYSIS 


Fig.  24. 


termined.  By  means  of  a  T-tube  this  standard  burette  is  made 
to  connect  with  the  burette  in  which  the  volume  of  air  to  be 
analyzed  is  measured  and  with  a  movable  reservoir 
of  mercury.  Both  burettes  are  firmly  fastened  to 
a  movable  iron  rod  and  the  zero  marks  accurately 
adjusted  to  the  same  level.  By  drawing  into  the 
measuring  burette  a  volume  of  moistened  air  greater 
than  that  required,  and  then,  by  raising  the  reser- 
voir of  mercury,  compressing  the  air  in  the  standard 
tube  to  the  100  cm.3  mark,  at  the  same  time  allow- 
ing the  excess  of  air  to  escape  from  the  measuring 
burette,  exactly  100  cm.3  of  air  under  the  standard 
conditions  of  temperature  and  pressure  may  be 
obtained.  To  facilitate  the  adjustment,  two  strips 
of  wood  are  fastened  to  the  rubber  connection  by 
means  of  screw  pinchcocks  in  such  a  way  that  by 
LEVELLER  closing  one  pinchcock  the  flow  of  mercury  from 
the  reservoir  is  shut  off,  while  by  tightening  the 
other  pinchcock  mercury  is  gradually  forced  out 
of  the  rubber. 

The  action  of  the  oxygen  upon  hydriodic  acid  is  effected  in  a 
300  cm.3  bulb  pipette,  both  ends  of  which  are  cut  off  short  and 
sealed  to  glass  stopcocks.  The  tube  from  one  of  the  stopcocks 
is  cut  off  short  after  being  tapered  and  constricted  so  as  to  hold 
a  rubber  connector  tightly,  while  the  tube  from  the  other  stop- 
cock is  left  sufficiently  long  to  reach  to  the  bottom  of  a 
500  cm.3  Erlenmeyer  beaker.  These  tubes  are  preferably 
of  about  3  mm.  bore,  since  for  the  several  connections 
all  air  may  be  expelled  from  tubes  of  this  size  by  dis- 
placement with  water.  The  bulb  is  filled  with  water  to 
expel  all  air.  The  water  is  then  displaced  by  pure 
carbon  dioxide  (prepared  as  described  below)  and  the 
flask  is  exhausted  by  connecting  it  with  a  large  bottle 
kept  vacuous  by  a  water  pump.  The  required  amounts 
of  potassium  iodide  solution,  hydrochloric  acid  and 
nitrogen  dioxide  are  introduced  in  the  order  named, 
after  which  the  measured  volume  of  air  is  gradually 
admitted  while  the  bulb  is  constantly  agitated  to  keep  the 
hydriodic  acid  continually  renewed  upon  the  surface  of  the 
bulb.  The  shaking  is  continued  for  a  minute  or  two  until 


Fig.  25. 


OXYGEN  357 

the  action  is  completed,  when  a  dilute  solution  of  potassium 
bicarbonate  is  admitted.  Upon  opening  the  lower  stopcock,  the 
pressure  of  the  carbon  dioxide  forces  the  liquid  from  the  bulb  into 
a  solution  of  the  bicarbonate  in  amount  sufficient,  as  previously 
determined,  to  neutralize  all  the  acid  taken.  Bicarbonate  is 
again  admitted  through  the  upper  stopcock,  and  after  neutrali- 
zation has  been  completed  the  bulb  may  be  washed  out  without 
any  danger  from  the  admission  of  air. 

All  the  water  employed,  both  for  the  solution  of  potassium 
iodide  and  for  the  various  connections,  must  be  free  from  oxygen. 
It  is  prepared  by  filling  a  three  liter  flask  with  distilled  water  and 
boiling  until  the  volume  of  the  liquid  is  reduced  about  one-third, 
when  the  flask  is  closed  by  a  doubly  perforated  rubber  stopper 
fitted  with  tubes  like  a  wash  bottle.  By  means  of  the  tube  which 
reaches  below  the  surface  of  the  water,  pure  carbon  dioxide  is 
passed  through  while  the  water  is  still  boiling.  Then  the  source 
of  heat  is  removed  and  the  escape  tube  is  closed  by  a  piece  of 
rubber  tubing  and  screw  pinchcock.  As  the  water  cools  it  is  well 
shaken  while  carbon  dioxide  is  admitted  from  the  generator,  and 
finally  pumped  in  under  considerable  pressure  by  the  little  hand 
pump  shown  in  Fig.  9.  From  this  supply  the  water  may  be  drawn 
as  needed  without  the  introduction  of  any  air,  the  escape  tube 
being  provided  with  a  rubber  tube  and  screw  pinchcock  and  a 
long,  slender  nozzle  which  may  be  inserted  into  the  tubes  of  the 
absorption  apparatus.  A  bottle  thus  charged  suffices  for  many 
determinations  and  requires  only  occasional  recharging  with  car- 
bon dioxide. 

The  potassium  iodide  solution  is  made  up  to  contain  I  grm. 
of  the  salt  in  30  cm.3  of  water,  and  is  contained  in  an  ordinary 
wide-mouthed  bottle,  fitted  as  a  wash  bottle,  and  graduated 
approximately  to  volumes  of  30  cm.3  —  the  amount  usually 
taken.  In  making  the  solution,  the  potassium  iodide  is  weighed 
into  the  bottle,  which  is  then  closed.  All  air  is  expelled  by  a 
current  of  carbon  dioxide,  the  desired  amount  of  water,  free  from 
oxygen,  is  drawn  in,  and  attachment  is  again  made  with  the 
carbon  dioxide  generator.  After  allowing  the  gas  to  pass  for 
several  minutes  the  exit  is  closed  and  the  gas  is  pumped  in  by  the 
little  hand  pump.  Inasmuch  as  this  solution  is  drawn  when  used 
into  an  exhausted  bulb,  the  bottle  may  be  emptied  without  expos- 
ing its  contents  to  the  air. 


358 


METHODS  IN  CHEMICAL  ANALYSIS 


Nitrogen  dioxide  is  generated  by  the  action  of  nitric  acid  upon 
granulated  shot  copper  in  a  Kipp  generator.  When  the  nitric 
acid  is  diluted  with  an  equal  volume  of  water  the  evolution  of 
the  gas  is  sufficiently  rapid  without  the  application  of  heat,  but 
contamination  by  the  higher  oxide  is  more  likely.  However, 
by  passing  the  gas,  as  it  issues  from  the  generator,  through  a  set 
of  Geissler  bulbs  containing  an  acidified  solution  of  potassium 
iodide  and  washing  with  potassium  iodide  solution,  perfectly 
purified  gas  is  obtained.  Theoretically,  only  a  small  amount  of 
the  nitrogen  dioxide  is  required  for  the  transference  of  the  oxygen 
to  the  hydriodic  acid,  but  when  too  little  is  taken  the  action  is 
very  slow.  On  the  other  hand,  too  large  an  amount  relieves  the 
vacuum  to  such  an  extent  as  to  interfere  with  the  introduction  of 
the  air.  A  little  device  to  measure  the  volume  of  gas  taken  is 
therefore  attached  to  the  generator.  It  consists  of  a  tube  filled 
with  water  and  roughly  graduated  for  every  five  cubic  centime- 
ters, so  attached  to  the  generator  that  the  gas  enters  with  displace- 
ment of  the  water  to  a  lower  bulb,  and  as  it  is  withdrawn  is  replaced 
by  water.  A  volume  of  fifteen  cubic  centimeters  of  the  gas  is  a 
convenient  and  satisfactory  amount  for  an  analysis. 

Carbon  dioxide  is  generated  in  a  Kipp  generator,  which  is 
charged  with  previously  boiled  acid  and  marble  and  a  little 
cuprous  chloride.  To  remove  traces  of  reducing  matter,  it  is 
first  passed  through  a  solution  of  iodine  and  then  washed  with 

potassium  iodide. 

Relation  of  Arsenic  to  Oxygen. 


w/io  As2O3. 
cm.8 

Oxygen  equivalent 
at  o°  and  760  mm. 

cm.3 

Correction  for 
o.oi  cm.3  «/io 
As203. 

37-0 

20.714 

0.005 

37-i 

20.770 

37-2 

20.826 

37-3 

20.882 

37-4 

20.938 

37-5 

20.994 

37-6 

21.050 

37-7 

21  .IO6 

37-8 

21  .162 

37-9 

21  .2l8 

38.0 

21.274 

For  the  titration  a  decinormal  solution  of  arsenious    oxide 
(4-95  grm-  to  the  liter)  was  employed,  one  cubic  centimeter 


OXYGEN 


359 


being  equal  to  0.55985  cm.3  of  oxygen  at  o°  and  760  mm.  when 
the  weight  of  a  liter  of  oxygen  at  o°  and  760  mm.  is  taken  as 
1.42895  grm.  When  the  volume  of  air  taken  is  100  cm.3  under 
standard  conditions  of  temperature  and  pressure,  as  obtained 
by  Lunge's  device,  the  preceding  table,  calculated  for  the  volume 
of  oxygen  equivalent  to  the  volume  of  arsenic  solution,  shows 
directly  the  percentage  of  oxygen  corresponding  to  the  reading 
of  the  burette.  The  correction  necessary  for  the  fraction  of  a 
tenth  of  a  cubic  centimeter  of  the  arsenic  solution  is  obtained 
with  sufficient  accuracy  by  simply  multiplying  the  figure  express- 
ing hundredths  of  a  cubic  centimeter  by  0.005. 

In  the  table  are  shown  results  obtained  upon  portions  of  a 
sample  of  air  either  measured  in  the  adaptation  of  Lunge's 
device  to  make  100  cm.3  under  standard  conditions  or  measured 
in  an  ordinary  gas  burette  and  calculated  to  standard  conditions. 

Oxygen  in  Air. 


Volume  of  air  reduced 
to  o°  and  760  mm. 

cm.8 

n/io  As?O3  required, 
cm.* 

Volume  of  oxygen 
found  at  o°  and 
760  mm. 

cm.3 

Per  cent  of  oxygen 
in  air. 

Measured  to  standard  conditions. 


IOO.OO 

37-44 

20.96 

20.96 

IOO.OO 

37-54 

21.01 

21  .OI 

IOO.OO 

37-50 

20.99 

20.99 

IOO.OO 

37-57 

21.03 

21.03 

IOO.OO 

37-47 

20.97 

20.97 

IOO.OO 

37-50 

20.99 

20.99 

Calculated  to  standard  conditions. 


91.18 

34-o6 

19.07 

20.91 

9i  73 

34-47  ' 

19.30 

21  .04 

90.84 

34-25 

19.17 

21  .  II 

90.60 

34-20 

19.  16 

21.13 

86.06 

32-55 

18.22 

21.17 

85.96 

32.40 

18.14 

21.10 

86.49 

32.53 

18.21 

21.  06 

87.85 

33-00 

18.47 

21.03 

44-17 

16.60 

9.29 

21  .04 

44.11 

16.  70 

9-35 

21  .  IO 

44-54 

16.80 

9.41 

21  .  12 

As  is  evident  from  the  table,  the  determinations  according  to 
this  method  are  not  reliable  beyond  0.05  per  cent,  but  for  prac- 
tical purposes  this  is  sufficiently  accurate. 


360  METHODS  IN  CHEMICAL  ANALYSIS 

Determination  of      ^  determination  of  oxygen  dissolved  in  water  may 
Dissolved  Oxy-   be  completed  in  about  ten  minutes  by  means  of  the 
apparatus  illustrated   by  the  accompanying   figure. 
The  apparatus  consists  of  a  flask  of  about  300  cm.3  capacity, 
into  the  bottom  of  which  is  sealed  a  stopcock  with  a  long  exit 
tube.     Upon  the  neck  is  cut  the  fiducial  circle  c  and 
above  the  stopcock  e  is  sealed  on,  as  shown.     The 
1 6  neck  of  the  flask  is  drawn  out  and  sealed  to  the  stop- 
cock d  and  the  bulb  a,  of  about  30  cm.3  capacity  blown 
in  it.     The  capacity  of  the  apparatus  between  stop- 
cock b,  and  the  fiducial  mark  c,  is  carefully  deter- 
mined. 

The  manipulation  for  the  determination  of  dissolved 
oxygen  is  as  follows :  The  flask  is  held  in  position  by 
a  clamp  fastened  to  a  movable  support.  Stopcock  b 
being  closed,  the  water  is  admitted  through  e  and 
Fig.  26.  the  air  allowed  to  escape  through  d  until  the  level 
of  water  is  that  indicated  by  the  line/.  (When  the  water  to 
be  examined  is  not  saturated  with  air,  the  flask  must  first  be 
filled  with  carbon  dioxide  and  the  water  introduced  by  replace- 
ment of  that  gas.)  With  d  closed,  sufficient  water  is  allowed 
to  escape  through  b  to  bring  the  surface  to  e,  which  is  then 
closed.  The  nitrogen  dioxide  generator  is  then  attached  to  d, 
and  by  opening  b  the  gas  is  allowed  to  replace  the  water  until 
the  meniscus  coincides  with  c,  when  d  is  closed  and  the  generator 
disconnected.  Two  cubic  centimeters  of  strong  hydrochloric  acid 
are  introduced  through  e  by  expelling  nitrogen  dioxide  through  d, 
in  which  a  drop  of  water  forms  an  effective  trap  to  prevent  the 
entrance  of  air.  Then  the  potassium  iodide  is  admitted  in  the 
same  way.  The  solution  of  iodide  for  this  purpose  is  free  from 
oxygen  and  contains  I  grm.  in  3  cm.3.  It  is  kept  under  pressure 
of  carbon  dioxide  in  the  bottle  previously  described,  and  by 
means  of  a  long  nozzle  may  be  conducted  to  the  bottom  of  eh 
and  thus  be  admitted  with  but  momentary  and  slight  contact 
with  the  air.  The  tube  eh  contains  approximately  3  cm.3.  With 
all  the  stopcocks  closed,  the  flask  is  inverted  several  times  and 
thoroughly  shaken,  and  the  ends  of  the  stopcocks  are  washed  out 
with  distilled  water.  After  again  placing  the  apparatus  in  its 
position,  enough  potassium  bicarbonate  solution  is  admitted 
through  e  to  expel  all  the  nitrogen  dioxide  through  d,  the  bulb  a 


OXYGEN 


361 


holding  sufficient  bicarbonate  to  neutralize  all  the  acid  taken. 
The  bicarbonate  being  heavier  quickly  diffuses  through  the 
contents  of  the  flask  and  neutralizes  the  acid;  d  and  e  are  kept 
closed  for  a  minute  with  b  open  so  as  to  allow  sufficient  of  the 
liquid  to  escape  into  a  beaker  containing  some  bicarbonate  to 
provide  space  for  the  carbon  dioxide  evolved.  Then  the  flask 
is  washed  out  into  the  beaker  and  its  contents  titrated  with 
standard  arsenite. 

The  bleaching,  by  the  aid  of  starch  for  the  final  reaction,  may 
be  accurately  read  to  a  single  drop  and  the  reading  verified 
by  adding  a  drop  of  n/io  iodine  solution  to  characteristic 
coloration. 

The  table  gives  the  results  of  a  series  of  determinations. 

Oxygen  in  Water. 


Volume  of  water  taken. 
cm.3 

Temperature. 
C°. 

As2O3  required. 
cm.3 

Volume  of  oxygen 
dissolved  in  1000  cm.3 
of  water  at  760  mm. 

314.63 

2O 

3-42 

6.04 

3I4-63 

2O 

3-45 

6.OQ 

3H.63 

2O 

3-4° 

6.00 

3U.63 

2O 

3-4i 

6.  02 

3H.63 

2O 

3-43 

6.05 

3I4-63 

20 

3-40 

6.00 

314-63 

20 

3.36 

5-93 

314.63 

2O 

3-40 

6.00 

314.63 

2O 

3-40 

6.00 

3M.63 

2O 

3-50 

6.18 

314.63 

2O 

3.38 

5-96 

3H.63 

2O 

3-40 

6.60 

The  mean  of  these  determinations  gives  6.022  cm.3  of  oxygen 
as  the  amount  dissolved  in  distilled  water  at  20°  and  760  mm., 
and  while  some  of  the  determinations  vary  considerably  from 
this  mean,  as  a  whole  they  are  fairly  accordant.  This  method, 
moreover,  is  applicable  to  carbonated  water. 

The  Estimation  of  Oxidizers  by  the  Gravimetric  Determination  of 
Iodine  Set  Free  in  Reaction. 

The  process  developed  by  Perkins  for  the  gravimetric  esti- 
mation of  iodine  by  absorption  in  metallic  silver  may  be  utilized 
for  the  determination  of  oxidizers.*    In  determining  available 
oxygen  by  this  process  a  definite  amount  of  the  oxidizer  is  added 
*  Claude  C.  Perkins,  Am.  Jour.  Sci.,  [4],  xxix,  339. 


362 


METHODS  IN  CHEMICAL  ANALYSIS 


to  a  solution  of  potassium  iodide  acidified  with  hydrochloric  acid 
and  the  mixture  is  shaken  with  a  weighed  amount  of  silver  in  an 
atmosphere  of  hydrogen.  The  value  of  the  oxidizer  is  then  cal- 
culated from  the  increase  in  weight  of  silver  which  represents 
the  amount  of  liberated  iodine.  The  table  shows  figures  obtained 
in  the  determination  of  potassium  permanganate,  hydrogen  dioxide* 
potassium  dichromate,  and  ferric  chloride  taken  in  standardized 
solutions.  In  all  of  the  results  the  error  is  well  within  reason- 
able experimental  variation. 

Iodine  Liberated  by  Oxidizer s. 


Silver  taken, 
grm. 

Oxygen  available 
in  oxidizer. 

grm. 

Iodine  found, 
grm. 

Oxygen  found, 
grm. 

Error, 
grm. 

Determination  of  potassium  permanganate. 


3.0000 

0.0123 

0.1956 

0.0123 

o.oooo 

3.0000 

0.0123 

0.1936 

O.OI22 

—  O.OOOI 

3.0100 

0.0123 

0.1964 

0.0124 

+0.0001 

3.0100 

0.0123 

0.1968 

O.OI24 

+O.OOOI 

4.0101 

0.0185 

0.2926 

O.Ol84 

—  O.OOOI 

4.0101 

0.0247 

0.3886 

0.0245 

—  O.OOO2 

Determination  of  hydrogen  peroxide. 


3.0000 

O.O2O2 

0.3200 

O.O2O2 

o.oooo 

3.0000 

O.O2O2 

0.3214 

o  .  0203 

+O.OOOI 

3.0100 

o  .  0404 

0.6427 

o  .  0405 

-f-o.oooi 

3.0000 

0.03II 

0-4937 

0.03H 

0.0000 

3.0000 

0.0322 

0.5128 

0.0323 

-j-o.oooi 

3.0000 

o  .  0606 

0.9590 

o  .  0604 

—  O.OOO2 

Determination  of  potassium  dichromate. 


3.0000 

0.0080 

0.1272 

0.0080 

o.oooo 

3.0000 

0.0160 

0.2552 

0.0161 

+O.OOOI 

3.0000 

O.02OI 

0.3141 

0.0198 

—0.0003 

3.0000 

o  .  0402 

o  .  6390 

o  .  0403 

+0.0001 

3.0000 

0.0160 

0.2552 

0.0161 

+O.OOOI 

3.0000 

0.0160 

0.2571 

0.0162 

+O.OOO2 

Determination  of  ferric  chloride. 


3.0000 

0.0218 

0.3470 

0.0219 

+O.OOOI 

3.0000 

o.  2018 

0.3476 

0.0219 

+O.OOOI 

3.0000 

0.0437 

0.6922 

0.0436 

—  O.OOOI 

3.0000 

0.0218 

0.3489 

O.O22O 

+0.0002 

3.0000 

0.0262 

0.4183 

O.O264 

+O.OOO2 

3.0000 

0.0218 

0.3516 

0.0222 

+0.0004 

SULPHUR 


363 


SULPHUR. 

The  Detection  of  Sulphides,   Sulphates,   Sulphites  and   Thiosul- 

phates  in  the  Presence  of  One  Another. 

A  method  by  Browning  and  Howe  *  presents  improved  pro- 
cedure upon  lines  suggested  by  the  method  of  R.  Grieg  Smith,  f 
for  the  detection  of  sulphides,  sulphates,  sulphites  and  thio- 
sulphates.  The  method  as  modified  may  be  thus  summarized: 
To  about  o.i  grm.  of  the  substance  to  be  analyzed  dissolved  in 
10  cm.3  of  water  or  more,  add  sodium,  potassium,  or  ammonium 
hydroxide  to  distinct  but  faintly  alkaline  reaction.  The  solution 
should  be  neutral  or  alkaline  rather  than  even  faintly  acid,  owing 
to  the  readiness  with  which  sulphur  separates. 

(I)  To  the  alkaline  solution  add  zinc  acetate  in  distinct  excess 
and  filter.     The  precipitate  is  treated  with  acid  and  a  test  made 
for  hydrogen  sulphide,  to  indicate  presence  or  absence  of  a  sul- 
phide. 

(II)  To  the  filtrate  add  acetic  acid,  a  few  drops  in  excess  of  the 
amount  necessary  to  neutralize,  and  barium  chloride.     A  pre- 
cipitate of  barium  sulphate  indicates  a  sulphate. 

(III)  Filter   through   a   double   filter.     To   the   filtrate   add 
iodine  until  the  solution  takes  on  a  permanent  yellow  tinge,  and 
then  bleach  with  stannous  chloride,  best  after  adding  a  few  drops 
of  hydrochloric  acid  to  prevent  the  precipitation  of  a  basic  salt 
of  tin.     A  precipitate  of  barium  sulphate  indicates  a  sulphite. 

(IV)  Filter,  add  bromine  water  in  faint  excess  to  the  filtrate* 
bleaching  again  with  stannous  chloride.     A  precipitate  of  barium 
sulphate  indicates  a  thiosulphate  originally  present. 

The  results  of  test  experiments  are  given  below. 

Tests  for  Sulphite  after  Iodine  Treatment  (III). 


K,SO, 
taken. 

Volume 
of  water. 

BaSO4  precipitated 
after  oxidation 

Remarks. 

grm. 

cm.s 

with  iodine. 

O.  I 

IO 

Very  abundant. 

Plainly  visible  before  adding  SnCU. 

O.OI 

IO 

Abundant. 

Plainly  visible  before  adding  SnC^. 

0.001 

10 

Distinct. 

More  distinct  after  adding  SnCl2. 

0.0005 

10 

Fair. 

Hardly  visible  before  adding  SnCl2* 

O.OOOI 

10 

Faint. 

Invisible  before  adding  SnCl2. 

*  Philip  E.  Browning  and  Ernest  Howe,  Am.  Jour.  Sci.,  [4],  vi,  317. 
t  Chem.  News,  Ixxii,  39. 


364 


METHODS  IN  CHEMICAL  ANALYSIS 


Test  for  Thiosulphate  After  Iodine  and  Bromine  Treatment  (III)  and  (IV). 


Na2SaO, 

taken. 

Volume 
of  water. 

BaSO4  precipi- 
tated by  action 

BaSO4  precipi- 
tated by  action 

Remarks. 

grin. 

cm.* 

of  iodine. 

of  bromine. 

O.I 

IO 

Faint. 

Abundant. 

(  No    sulphur  separated    in 
)      i  minute. 

O.OI 

10 

None. 

Abundant. 

i  No    sulphur    separated  in 
(      several  minutes. 

O.OOI 

IO 

None. 

Distinct. 

No  sulphur. 

0.0005 

IO 

None. 

Faint. 

j  No   sulphur;    SnCl2   neces- 
J      sary. 

0.0001 

10 

None. 

Very  faint. 

(  No    sulphur;    SnCl2  neces- 
(      sary. 

Test  for  Sulphate  (III)  and  Thiosulphate  (IV),  After  Removal  of  Sulphide  (I) 

and  Sulphate  (II). 


K2SO,  taken. 

Na2S2O3  taken. 

BaSO«  precipitated  after 
oxidation  with  iodine. 

BaSO4  precipitated  after 
oxidation  with  bromine. 

grm. 

grm. 

O.I 

O.OI 

Abundant. 

Good. 

O.I 

O.OOI 

Abundant. 

Distinct. 

O.OI 

O.  I 

Good. 

Abundant. 

O.OOI 

O.  I 

Faint. 

Abundant- 

O.OOI 

O.OOI 

Fair. 

Fair. 

The  lodometric  Determination  of  Thio sulphates. 

In  an  investigation  of  the  reaction  between  iodine  and  sodium 
thiosulphate,  Pickering*  has  shown  that  more  iodine  is  required 
to  oxidize  the  thiosulphate  as  the  proportion  of  hydrochloric 
acid  increases.  This  effect  is  by  him  ascribed  to  the  formation 
of  sulphate,  apparently  by  the  increased  activity  of  the  iodine; 
but  the  more  rational  explanation  is  that,  although  some  sulphate 
Is  ultimately  formed,  the  thiosulphate  is  first  partially  decomposed 
Into  free  sulphur  and  sulphur  dioxide.  Finkener  \  and  Mohr  %. 
also  mention  the  decomposing  effect  of  free  acid  upon  sodium 
thiosulphate. 

Norton  §  has  studied  carefully  the  effect  of  acidity  in  the 
Iodine  titration  of  thiosulphate,  with  results  given  in  the  follow- 
ing account. 

*  Jour.  Chem.  Soc.,  vol.  xxxvii,  pages  135. 

f  Anal.  Chem.,  6  Aufl.,  pages  620. 

J  Titrirmethode,  6  Aufl.,  pages  279. 

§  John  T.  Norton,  Jr.,  Am.  Jour.  Sci.,  [4],  vii,  287. 


SULPHUR 


365 


The  sodium  thiosulphate  used  in  the  following  experiments, 
taken  in  nearly  decinormal  solution,  was  standardized  by  run- 
ning it  into  an  approximately  decinormal  solution  of  iodine,  the 
value  of  which  had  been  determined  by  comparison  with  deci- 
normal arsenious  acid  made  from  carefully  resublimed  arsenious 
oxide.  In  the  experiments  recorded  below  the  solutions  were 
stirred  continuously  while  the  thiosulphate  ran  into  the  acidified 
liquid.  The  results  show  clearly  that  the  amount  of  iodine  used 
in  the  titration  of  a  fixed  amount  of  sodium  thiosulphate  is 
dependent  to  a  very  marked  degree  upon  the  concentration  of 
the  thiosulphate  in  presence  of  the  acid,  the  concentration  of 
the  acid,  the  time  and  the  temperature. 


Varying  Concentrations  of  Thiosulphate  and  Acid;    Temperature  o°- 
tions  made  Rapidly. 


Titra- 


Volume  of 
liquid  at 
beginning  of 
titration. 

cm.8 

Na2S20, 

approximately 
«/io  taken. 

cm.* 

Volume  of  w/io  iodine  used  in  titration. 

HCl=none. 
cm.1 

=  i  cm.* 
cm.3 

=5  cm.' 
cm.1 

=  10  cm.' 
cm.* 

100 

3° 

30-25! 

T30.75 

30.76 

31-2 

200 

30 

3°-22  Imean- 

I  30.21 

30.56 

31-4 

300 

3° 

30-2or30  22 

•<  30.22 

3L03 

30.9 

400 

30 

30.21    6°- 

30.20 

30.20 

30.55 

500 

30 

30.20) 

130.20 

30.21 

30.55 

IOO 
200 
300 

25        • 
25 
25 

25-29! 

sifer 

f25-32 

25.34 
H  25.41 

25.98 
25.40 
25.38 

25.70 
25-45 
25.83 

400 

25 

25.27  125-27 

1  25.24 

25.30 

25-63 

500 

25 

25.22) 

125-23 

25.40 

25.30 

IOO 

2O 

20.15! 

r  20.17 

20.33 

20.23 

2OO 

2O 

20-20!mean= 

1  20.13 

20.27 

20.23 

300 

20 

2°-2If2o  15 

<j  20.15 

2O.  2O 

20.17 

400 

20 

20.20 

20.  10 

20.27 

20.07 

500 

2O 

2O.  IOJ 

(^20.10 

20.  17 

20.13 

Varying  Concentrations  of  Thiosulphate  and  Times;    Temperature  o°-5°. 


Volume  of  the 
liquid  at 
beginning  of 
titration. 

HC1 
sp.  gr.  I.I2 
present. 

Na2S203 
approximately 
w/io  taken. 

Volume  of  w/io  iodine  used  in  titration 
after  standing. 

5  minutes. 

10  minutes. 

15  minutes. 

cm.* 

cm.* 

cm.* 

cm.* 

cm.* 

cm.* 

2OO 
2OO 
2OO 

10 
IO 
IO 

30 
25 
20 

30.80 
25.50 
20.30 

31.30 
26.00 
20.70 

32.32 
26.30 
20.68 

366 


METHODS  IN  CHEMICAL  ANALYSIS 


Varying  Temperatures. 


Volume  of  liquid 
at  beginning  of 
titration. 

HC1 
sp.  gr.  1.12 

taken. 

Temperature. 

Na2S203 
approximately 
w/io  taken. 

Volume  of  w/io 
iodine  used  in 
titrations  at  differ- 
ent temperatures. 

cm.» 

cm.J 

cm.8 

cm.* 

400 

IO 

6° 

25 

23.52 

400 

IO 

22° 

25 

23  73 

400 

IO 

34o 

25 

24-35 

400 

IO 

42° 

25 

24-5 

400 

10 

54° 

25 

25 

400 

IO 

64° 

25 

26.1 

From  these  results  it  is  plain  that  the  conditions  under  which 
considerable  amounts  of  sodium  thiosulphate  are  titrated  in 
presence  of  hydrochloric  acid  must  be  carefully  guarded  when 
accuracy  is  a  consideration.  The  amount  of  acid  should  be 
restricted,  the  temperature  should  be  reduced  as  nearly  to  o° 
as  possible,  and  titration  by  the  iodine  should  be  made  promptly. 
So  long  as  the  thiosulphate  present  does  not  exceed  20  cm.3  of 
the  n/io  solution  in  a  volume  of  200  cm.3,  rapid  titration  in  cold 
solution  proceeds  with  fair  regularity  in  presence  of  hydro- 
chloric acid  up  to  10  cm.3  of  the  acid  of  sp.  gr.  1.12. 

Fortunately,  in  most  analytical  processes  involving  the  use 
of  the  thiosulphate  as  a  reagent  it  is  possible  to  add  that  reagent 
from  the  burette  to  the  solution  to  be  acted  upon,  so  that  it 
is  destroyed  normally  as  fast  as  it  is  introduced  and  the  danger 
of  interaction  with  the  acid  does  not  occur. 


The  lodometric  Determination  of  Sulphites  in  Alkaline  Solution. 

According  to  Bunsen,  Dupasquier's  method  of  oxidizing 
sulphurous  acid  by  iodine  in  an  acid  solution,  proceeds  to  com- 
pletion when  the  concentration  of  the  sulphur  dioxide  does  not 
exceed  0.5  per  cent  of  the  solution.  When,  on  the  other  hand, 
the  proportion  of  sulphur  dioxide  exceeds  this  value,  there  is  a 
secondary  reaction,  which,  according  to  Volhard,  involves  the 
reduction  of  the  sulphur  dioxide  by  the  hydriodic  acid  produced. 
The  reaction  proceeds  normally  in  dilute  solutions  according  to 
the  equation 

2SO2  +  2  I2  +  4H2O  =  4HI  +  2  H2S04. 


SULPHUR  367 

In  solutions  too  concentrated,  however,  the  secondary  reaction, 
S02  +  4  HI  =  2  I2  +  2  H20  +  S, 

takes  place,  as  Volhard  has  shown,*  and  vitiates  the  indications. 
The  difficulty  may  be  obviated,  however,  as  has  been  shown  by 
Volhard, t  if  the  solution  of  the  sulphurous  acid  or  a  sulphite  is 
run  with  stirring  into  a  solution  of  iodine  in  potassium  iodide, 
acidified  with  hydrochloric  acid,  to  the  bleaching  of  color,  using 
starch  as  an  indicator.  Volhard 's  method  is  accurate  and  reli- 
able, but  it  involves  the  inconvenience  of  making  up  every 
solution  to  be  examined  accurately  to  a  standard  volume  of 
which  portions  are  to  be  drawn  from  a  burette  and  made  to 
react  with  definite  amounts  of  a  standardized  solution  of  iodine. 
It  rests  upon  the  facts  that  the  oxidation  of  sulphite  is  brought 
about  in  the  acidified  solution  and  that  no  more  than  a  small 
proportion  of  hydriodic  acid  is  present  at  the  point  at  which  the 
oxidation  of  sulphur  dioxide  takes  place. 

To  avoid  the  inconvenience  of  the  Volhard  method  it  has 
been  proposed  by  Rupp  t  to  bring  about  the  oxidation  of  sul- 
phites by  treatment  with  an  excess  of  standardized  iodine  in  a 
solution  made  alkaline  by  acid  sodium  carbonate,  and  then, 
after  fifteen  minutes,  to  titrate  the  excess  of  iodine  by  sodium 
thiosulphate.  This  procedure,  however,  as  has  been  shown  by 
Ruff  and  Jaroch  §  and  by  Ashley J|  is  faulty  in  principle  and 
practice,  and  gives  correct  results  only  by. a  chance  balancing 
of  opposing  errors.  It  has  been  further  pointed  out  by  Ashley  ** 
that  it  is  possible  to  overcome  the  difficulties  by  treating  the 
alkaline  mixture  with  acid  before  attempting  to  titrate  by 
sodium  thiosulphate  the  excess  of  iodine.  According  to  Ashley's 
procedure,  the  practical  estimation  of  sulphurous  acid  or  of  a 
soluble  sulphite  may  be  accomplished  with  a  reasonable  degree 
of  accuracy  by  adding  to  the  solution  of  the  substance,  not 
exceeding  100  cm.3  in  volume  and  containing  a  gram  of  acid 
sodium  carbonate,  at  least  twice  as  much  iodine  as  is  theoreti- 
cally necessary  to  effect  oxidation,  acidifying  cautiously  with 

*  Ann.  Chem.  242,  93. 
f  Loc.  cit. 

t  Ber.  Dtsch.  chem.  Ges.,  xxxv,  3694. 
§  Ber.  Dtsch.  chem.  Ges.,  xxxviii,  409. 
II  Am.  Jour.  Sci.,  [4],  xix,  237. 
**  R.  Harmon  Ashley,  Am.  Jour.  Sci.,  [4],  xx,  1.3. 


368 


METHODS   IN  CHEMICAL  ANALYSIS 


hydrochloric  acid,  and  determining  with  standard  sodium  thio- 
sulphate  the  excess  of  iodine  remaining  in  the  acidified  solution. 
Results  of  this  procedure  are  given  in  the  table. 


Determination  of  Sulphite. 


Iodine 
value  of 
SO2  taken. 

grin. 

Iodine 
taken. 

grm. 

Iodine 
value  of 
Na2S203 
used. 

grm. 

Error. 

Excess  of 
HC1  [i  :  3.] 

cm.3 

Volume  at 
titration. 

cm.3 

In  terms 
of  iodine. 

grm. 

In  terms 
of  SO2. 

grm. 

0.1143 

0.3143 

0.1990 

+O.OOIO 

+0.0003 

5-0 

125 

0.1143 

0.3H3 

0.1982 

+O.OOI8 

+0.0004 

5-o 

125 

0.1143 

0.3H3 

o  1992 

-1-0.0008 

+0.0002 

5-o 

I25 

0.1143 

0.3H3 

0.1986 

+0.0014 

+0.0003 

5-° 

I25 

o.  1482 

0.3187 

O  .  I  708 

—  0.0003 

—  O.OOOI 

7-5 

I25 

0.1576 

0.3187 

0.1586 

+0.0025 

+O.OOO6 

7-5 

I25 

0.1576 

0.3187 

0.1643 

—  0.0032 

—  O.OOOS 

7-5 

125 

0.1576 

0.3187 

0.1598 

+0.0013 

+0.0003 

7-5 

125 

0.1576 

0.3187 

0.1606 

+0.0005 

+O.OOOI 

7-5 

I25 

0.1576 

0.3187 

o.  1602 

+o  .  0009 

+O.OO02 

7-5 

I25 

0.1576 

0.3187 

0.1622 

—  o.oon 

—  0.0003 

7-5 

125 

0.1560 

0.3195 

0.1660 

—0.0025 

—  O.OOO6 

7-5 

125 

0.1992 

o  .  4460 

o.  2482 

—0.0014 

—  0.0003 

7-5 

125 

0.1915 

0-3825 

0.1919 

+0.0009 

—  O.OOO2 

7-5 

125 

0.2056 

0.3771 

0.1701 

+0.0014 

+0.0003 

7-5 

I25 

0.2056 

0.3771 

0.1697 

+0.0018 

+  O.OOO4 

7-5 

125 

0.2056 

0.3771 

0.1707 

+0.0008 

+O.OO02 

7-5 

125 

0.2056 

0.3771 

0.1709 

+0.0006 

+  O.OOO2 

7-5 

125 

[0.2131 

0.4470 

0.2412 

-0.0073  • 

—  O.OOlS] 

7-5 

125 

0.2354 

0.3825 

0.1490 

—0.0019 

—  O.OOO5 

7-5 

125 

0.2597 

0.4463 

0.1869 

—0.0003 

—  O.OOOI 

7-5 

I25 

0.2638 

0.4463 

0.1847 

—  0.0022 

—0.0005 

7-5 

125 

0.2908 

0.6375 

0-3505 

—  0.0038 

—  0.0009 

7-5 

125 

0.3187 

0.4463 

o.  1326 

+O.O05O 

+O.OOI2 

7-5 

I25 

0-3395 

0.6275 

0.2842 

+0.0038 

+  O.0009 

7-5 

125 

0-3395 

0.6275 

0.2852 

+0.0028 

+O.O007 

7-5 

125 

0-3395 

0.6275 

o  .  2844 

+0.0036 

+O.OO09 

7-5 

I25 

0-3395 

0.6275 

0.2855 

+O.OO25 

+0  .  0006 

7-5 

125 

It  is  probable  that  the  formation  of  a  small  amount  of  dithio- 
nate  instead  of  sulphate  is  the  occasion  of  the  deficient  expenditure 
of  iodine  noced  when  the  concentration  of  this  element  is  low, 
and  that  the  dithionate  is  not  formed  appreciably  when  the 
iodine  concentration  is  high.  The  dithionate  once  formed  is 
but  slowly  attacked  by  iodine,  and  that  is  apparently  the 
reason  why  long  standing  of  the  mixtures  containing  a  small 
proportion  of  iodine  does  not  result  in  complete  oxidation  of 
the  sulphite  to  sulphate. 


SULPHUR 


369 


The  Determination  of  Dithionic  Acid  and  Dithionates. 

That  dithionic  acid  may  be  best  liberated  from  combination 
by  the  action  of  sulphuric  acid,  rather  than  hydrochloric  acid, 
has  been  shown  by  Ashley.*  In  studying  these  reactions, 
Ashley  treated  a  specially  prepared  barium  dithionate  with  the 
acids  in  the  distillation  apparatus  previously  described  and 
figured,!  and  estimated  the  resulting  sulphur  dioxide  by  absorb- 
ing it  in  standard  iodine  the  excess  of  which  was  determined 
by  titration  with  thiosulphate.  The  successful  procedure  is  as 
follows:  A  weighed  amount  of  dithionate  is  introduced  into 
the  Voit  flask  and  there  dissolved  in  water.  Sulphuric  acid 
is  run  in  through  the  separating  funnel  and  the  mixture  then 
boiled,  the  sulphur  dioxide  being  collected  in  the  Drexel  receiver 
charged  with  a  measured  amount  of  standard  iodine  in  potas- 
sium iodide  and  trapped  with  potassium  iodide.  A  slow  current 
of  carbon  dioxide  is  driven  through  the  system  to  sweep  the 
sulphur  dioxide  into  the  iodine  and  to  prevent  any  sucking 
back.  When  boiling  has  been  carried  so  far  that  fumes  of 
sulphuric  acid  begin  to  appear  the  operation  is  stopped,  and 
the  excess  of  iodine  remaining  is  determine^  by  means  of  sodium 
thiosulphate,  starch  iodide  being  used  as  an  indicator.  Results 
of  experiments  carried  out  in  this  manner  with  barium  dithionate 
are  given  in  the  table. 

Decomposition  of  Barium  Dithionate  by  Boiling  with  Sulphuric  Acid. 


S206 

ta.<en. 

I  value 
of  S205 
taken. 

I  taken. 

I  value  of 
of  Na2S203 
required. 

S20S 

found. 

Errors  in 
I. 

Errors  in 
S205. 

Time. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

min. 

0.1039 

0.2310 

0-5759 

0-3435 

0.1045 

+0.0014 

+0.0006 

20 

o.  1046 

o  .  23  26 

0.5708 

0.3372 

0.1051 

+O.OOIO 

+0.0005 

28 

0.1039 

o.  2311 

0.5740 

0-3435 

0.1037 

—  O.OOO6 

—  0.0002 

34 

0.1033 

0.2297 

0.5701 

0.3387 

o.  1041 

+0.0017 

+0.0008 

45 

o.  1721 

0.3827 

0.5712 

0.1876 

o.  1726 

+0.0009 

+o  .  0005 

35 

0.1719 

0.3820 

0.5702 

0.1894 

0.1713 

—  O.OOI2 

—  O.OOO6 

50 

0.1726 

0.3838 

0-5734 

0.1898 

0.1726 

—  O.OOO2 

o.oooo 

12 

0.1724 

0.3832 

0.5727 

0.1885 

0.1728 

+O.OOIO 

+0.0004 

10 

o.  1721 

0.3826 

0.3727 

0.1886 

o.  1728 

+0.0015 

+0.0007 

10 

o  0692 

0.1539 

0.3130 

0.1599 

0.0689 

—  O.OOOS 

—0.0003 

12 

0.0350 

0.0777 

0.3109 

0.2323 

0.0354 

+o  .  0009 

+o  .  0004 

4 

0.2061 

0.4582 

0.6205 

0.1632 

0.2057 

—  O.OOO9 

—0.0004 

15 

o.  2402 

0.5340 

0.6215 

0.0862 

0.2408 

+0.0013 

+0  .  0006 

15 

*  R.  Harmon  Ashley,  Am.  Jour.  Sci.,  [4],  xxii.  259. 
t  See  Fig.  3,  page  4. 


370  METHODS  IN  CHEMICAL  ANALYSIS 

Similar  experiments  in  which  hydrochloric  acid  was  substituted 
for  sulphuric  acid  were  not  satisfactory,  chiefly  because  the 
decomposition  of  the  dithionate  was  not  complete  even  after 
long  and  repeated  treatments,  and  the  large  amount  of  the  hydro- 
chloric acid  carried  to  the  receiver  tends,  as  is  well  known,  to 
render  titration  of  the  residual  iodine  by  thiosulphate  less  exact, 
and  the  starch  iodide  less  delicate  as  an  indicator. 

There  are  reasons  why  sulphuric  acid  should  work  better  than 
hydrochloric  acid  in  this  process: 

First,  when  sulphuric  acid  is  added  to  the  solution  of  barium 
dithionate,  barium  sulphate  is  precipitated  and  dithionic  acid 
is  left  in  free  condition,  this  reaction  proceeding  at  once  to 
completion  because  the  barium  sulphate  formed  is  removed 
from  the  system.  When  hydrochloric  acid  is  used  the  dithionic 
acid  is  completely  liberated  only  by  a  gradual  change  in  the 
conditions  of  equilibrium. 

Second,  when  the  solution  containing  sulphuric  acid  is  boiled, 
the  water  is  driven  off,  the  concentration  of  the  solution  increases 
and  the  temperature  of  the  sulphuric  acid  reaches  a  high  point. 
Under  such  conditions  the  decomposition  of  the  dithionic  acid  is 
rapid  and  complete,  'the  time  being  dependent  upon  the  original 
dilution  of  the  solution. 

Third,  no  appreciable  amount  of  acid  distils  over  from  the 
Voit  flask  into  the  receiver  containing  the  iodine,  to  interfere 
with  the  back  titration  with  sodium  thiosulphate,  the  only  acid 
present  being  that  produced  by  the  oxidation  of  the  sulphur 
dioxide.  Under  these  conditions  the  starch  indicator  acts 
sharply,  which  is  not  the  case  when  hydrochloric  acid  is  used. 

The  Determination  of  Per  sulphates. 

Arsenate-iodide  The  estimation  of  persulphates  is  accomplished 
Method.  by  Peters  and  Moody  *  by  procedure  similar  to  that 

of  Gooch  and  Smith  f  for  the  estimation  of  chlorates.  Accord- 
ing to  this  process  a  mixture  containing  the  persulphate,  potas- 
sium iodide,  2  grm.  to  3  grm.  of  hydrogen  potassium  arsenate, 
20  cm.3  of  [i  :  i]  sulphuric  acid,  and  water  enough  to  make  a 
total  volume  of  about  100  cm.3,  is  boiled  in  a  trapped  Erlen- 

*  Charles  A.  Peters  and  Seth  E.  Moody,  Am.  Jour.  Sci.,  [4],  xii,  367. 
t  See  page  463. 


SULPHUR 


371 


meyer  beaker  *  until  the  volume  is  diminished  to  35  cm.3,  when 
the  solution  is  made  alkaline  with  potassium  bicarbonate  and 
the  arsenite  present  is  estimated  by  standard  iodine  in  presence 
of  starch.  The  results  obtained  are  in  close  agreement,  as 
shown  by  the  record  below: 


A  r senate-Iodide  Method. 


Ammonium 
persulphate 
solution. 

KI 

present  . 

Iodine 
required  for 
oxidation  of 
of  arsenite. 

Iodine 
liberated  by 
persulphate. 

(NH4)2S208 
equivalent 
to  iodine 
liberated. 

(NH<)2S208 
average. 

cm.* 

gnu. 

grm. 

grm. 

grm. 

grm. 

12-5 

12.5 

I2-S 
12-5 

0.1875 
0.1875 
0.1875 
0.1875 

0.0514 
0.0522 
0.0514 
0.0516 

0.1361 
0.1352 
0.1361 
0.1.359 

0.1225    ^ 
0.1217     I 
0.1225     I 
O.I222     J 

0.1222 

Making  use  of  the  arsenate-iodide  method  as  a  control,  Peters 
and  Moody  have  brought  into  comparison  the  results  of  other 
and  simpler  methods,  viz.,  the  methods  of  LeBlanc  and  Eckardt, 
Griitzner,  Mondolfo  and  Namias. 

Method  of  According  to  LeBlanc  and  Eckardt,  f  when  a  mix- 

LeBianc  and  ture  containing  a  persulphate,  a  sufficient  excess  of 
ferrous  salt,  and  sulphuric  acid  is  heated  at  6o°-8o°, 
or  allowed  to  stand  ten  or  twelve  hours,  the  persulphate  is  re- 
duced, and  the  amount  of  ferrous  salt  oxidized  is  the  measure 
of  the  amount  of  persulphate  originally  present  in  solution. 
The  experiments  of  Peters  and  Moody  show  results  agreeing 
quite  closely  with  one  another  and  in  fair  agreement  with  the 
average  of  the  arsenate-iodide  method. 

Experiments  made  in  blank  to  discover  the  amount  of  ferrous 
salt  oxidized  in  eleven  hours  by  other  agencies  than  the  persul- 
phate showed  an  amount  of  oxidation  in  that  time  too  small  to 
be  detected ;  but,  when  the  solution  was  allowed  to  stand  thirty- 
six  hours,  a  slight  oxidation  was  noticed,  which,  if  calculated  as 
persulphate,  would  be  equivalent  to  0.0006  grm.  The  correction 
for  eleven  hours'  standing,  upon  this  basis,  therefore,  would  be 
0.0002  grm.,  which  is  obviously  so  small  that  it  may  be  disre- 
garded. This  absence  of  any  significant  oxidation  of  the  ferrous 

*  See  Fig.  6,  page  6. 

t  Zeit.  Elektrochem.,  5,  355-7- 


372 


METHODS  IN  CHEMICAL  ANALYSIS 


salt  was  undoubtedly  due  to  the  fact  that  the  solution  of  am- 
monio-ferrous  sulphate  had  been  standing  some  time  before 
standardizing  and  was  devoid  of  dissolved  oxygen. 

Ferrous  Sulphate  Method. 
Volume  of  liquid,  100  cm.3 


Ammonium 
persulphate, 
10  grm.  to 
liter. 

cm.» 

(NH4)2Fe 
(SO4)2.6H2O 
solution, 
20  grm.  to 
liter. 

cm.* 

KMn04 
approx.  w/io, 
to  oxidize  the 
excess  of 
ferrous  salt. 

cm.8 

(NH4)2S208 
calculated 
from  ferrous 
salt  oxidized. 

grm. 

(NH4)2S208 
average. 

grm. 

Variation 
from  results 
80°  of  arsenate- 
iodide  method. 

grm. 

Heated  10  minutes  at  6o°-8o°. 


12.  S 

25 

2.  II 

0.1216  ") 

I2.S 
12-5 
12-5 

25 
25 
25 

2.18 
2.  II 
2.  II 

O.I2OQ   1 

0.1216  I 
0.1216  j 

0.1217 

—0.0005 

12-5 

50 

15.00 

0.1218  | 

25.0 

50 

4.21 

0.2426  J 

Stood  ii  hours  at  2i°-25°. 


12.5 
12.5 
12.5 
12.5 

26 
26 
40 
40 

1.89 

1.  80 

8-75 
8-75 

0.1214  "I 
0.1224  I 

O.I22I   f 
O.I22I  J 

O.I22O 

—  0.0002 

Carbon  dioxide  above  the  liquid:   stood  n  hours  at  2i°-25°. 


12.5 
12.5 

40 
40 

8-75 
8.78 

O.I22I   } 

0.1218  ) 

O.I2IQ 

—  O.OO03 

Method  of  Grutzner  *  has  stated  that  arsenious  oxide  is  well 

Grutzner.  suited  to  the  determination  of  the  oxidizing   power 

of  persulphates,  the  reaction  being  hastened  by  heat  and  the 
presence  of  an  alkali  hydroxide. 

In  testing  the  value  of  the  persulphate  solution  according  to 
the  method  of  Grutzner,  the  mixture  containing  the  persulphate, 
standard  arsenite  solution,  and  sodium  hydroxide,  was  raised  to 
the  boiling  point,  cooled,  faintly  acidified  with  sulphuric  acid, 
then  made  alkaline  with  potassium  bicarbonate,  and,  after  adding 
a  large  excess  of  the  last,  the  arsenite  remaining  unoxidized  was 
*  Chem.  Centralblatt,  1900,  i,  p.  835. 


SULPHUR 


373 


titrated  by  iodine.  Experiments  made  in  blank,  —  that  is, 
experiments  in  which  the  mixture  of  arsenite  and  alkaline  hy- 
droxide was  raised  to  the  boiling  point,  cooled,  made  acid  with 
sulphuric  acid,  then  alkaline  with  potassium  bicarbonate,  and 
titrated  with  iodine,  —  showed  a  deficiency  of  arsenious  acid  as 
compared  with  the  experiments,  otherwise  similarly  conducted, 
in  which  the  mixture  was  not  subjected  to  heat.  The  effects  of 
various  specimens  of  sodium  and  potassium  hydroxides  were 
tested  in  the  process  —  sodium  and  potassium  hydroxides  made 
by  the  alcohol  process,  potassium  hydroxide  by  the  barium 
hydroxide  process,  sodium  hydroxide  especially  prepared  from 
sodium  carbonate  and  calcium  hydroxide  —  as  well  as  a  specimen 
of  potassium  carbonate.  With  all  excepting  the  potassium  car- 
bonate more  or  less  oxidation  was  observed,  sometimes  trifling, 
sometimes  considerable.  The  same  specimens  of  potassium  and 
sodium  hydroxides  and  potassium  carbonate  were  tested  in  the 
absence  of  the  arsenite  solution,  by  dissolving  them  in  water, 
heating  the  solution  to  the  boiling  point,  cooling,  adding  sul- 
phuric acid  in  faint  excess  and  then  potassium  bicarbonate, 
and  titrating  with  iodine  to  a  color  without  starch.  In  all 
cases,  save  that  of  the  potassium  carbonate,  the  color  of  the 
first  drop  of  iodine  solution  added  was  destroyed,  and  from 
0.06  cm.3  to  0.19  cm.3  were  necessary  to  bring  the  permanent 
iodine  color.  Potassium  carbonate,  however,  proves  to  be  in- 
efficient as  a  substitute  for  an  alkali  hydroxide.  By  using 
definite  amounts  of  the  hydroxide  and  introducing  a  correc- 
tion for  the  oxidation  determined  by  the  blank  tests  the  real 
oxidizing  effect  of  the  persulphate  upon  the  arsenite  may  be 
deduced.  The  table  contains  the  values  found  by  Griitzner's 
process,  both  corrected  and  uncorrected. 

Arsenious  Oxide  Method. 


Ammonium 
persulphate 
solution 
taken. 

Arsenite 
solution. 

NaOH  by 
Ca02H2. 

Calculated 
from  iodine 
actually  used. 

(NH4)2S208 
corrected  for 
oxidation  in 
blank. 

(NH4>2S208 
corrected, 
average. 

cm.3 

cm.3 

gnu. 

grm. 

grin. 

gnu. 

12-5 

25 

2 

0.1223 

0.1218  "] 

12.5 

25 

2 

0.1223 

0.1218    1 

O.  I2IQ 

12.5 

25 

2 

0.1225 

O.I22O      j 

12.5 

25 

2 

O.I22Q 

0.1220    J 

374  METHODS  IN  CHEMICAL  ANALYSIS 

Method  of  Marshall  *  is  authority  for  the  statement  that  per- 

Mondoifo.  sulphates  liberate  iodine  from  potassium  iodide,  and 
that  the  action  is  hastened  by  heat  and  affected  little  by  the 
addition  of  dilute  sulphuric  acid.  Upon  this  reaction  Mondolfo  f 
has  based  a  method  for  the  estimation  of  persulphates,  which 
consists  in  heating  a  mixture  containing  the  persulphate  and 
potassium  iodide  in  a  stoppered  bottle  for  ten  minutes  at 
60°  to  80°,  and  titrating  by  thiosulphate  the  iodine  set  free. 
According  to  Peters  and  Moody,  J  however,  the  reduction  of 
persulphate  is  incomplete  even  when  the  volume  of  the  mixture  is 
restricted  to  25  cm.3,  and  the  digestion  continued  thirty  minutes 
with  the  addition  of  small  amounts  of  sulphuric  acid.  Under  the 
best  conditions  the  amount  of  persulphate  found  was  0.1207  grm- 
in  an  amount  of  solution  for  which  the  arsenate-iodide  process 
indicated  0.1222  grm.,  and  the  LeBlanc  and  Eckardt  process 
0.1218  grm.  in  the  average. 

Method  of  The  process  proposed  by  Namias,  without  knowl- 

Namias.  e(jge    of    Mondolfo's    method  —  differs    from    the 

method  of  Mondolfo  in  the  particular  that  the  reaction  is  carried 
out  at  the  ordinary  temperature.  The  mixture  of  persulphate 
and  potassium  iodide  in  solution  is  allowed  to  stand  eleven  hours 
in  a  stoppered  bottle  and  the  iodine  set  free  is  titrated  with  thio- 
sulphate. In  experiments  conducted  by  Peters  and  Moody  § 
the  reduction  of  the  persulphate  under  the  conditions  described 
was  incomplete,  the  color  of  iodine  returning  upon  longer 
standing  after  the  titration.  Under  the  best  conditions  the 
value  of  0.1208  grm.  was  found  for  the  persulphate  contained  in 
an  amount  of  solution  for  which  the  arsenate-iodide  method  indi- 
cated o.i 222  grm.  and  the  LeBlanc  and  Eckardt  process  0.1218 
grm.  in  the  average. 

Comparison  To  compare  the  values  obtained  for  the  persul- 

of  Methods.  phate  solution  the  averages  of  the  results  obtained 
by  the  different  methods,  together  with  the  average  of  all  the 
experiments,  are  given  in  the  table. 

The  process  of  Mondolfo  and  the  process  of  Namias,  both  of 
which  involve  the  liberation  of  iodine  from  potassium  iodide  and 

*  Jour.  Chem.  Soc.  59,  771. 
f  Chem.  Ztg.,  23,  699. 
t  Loc.  cit. 
§  Loc.  cit.        • 


SELENIUM 


375 


the  titration  of  that  iodine  by  thiosulphate,  give  results  which 
are  practically  identical  and  lower*  than  those  obtained  by  the 
other  three  methods. 


Process. 

Number  of 
experiments. 

Average  of 
results.  ' 

grm. 

Mondolfo  

6 

O.I2O7 

Namias 

8 

o  1208 

LeBlanc  and  Eckardt 

12 

o  1217 

Griitzner  (corrected) 

4 

O.  I2IQ 

Arsenate-iodide  method  ... 

4 

0.  1222 

The  process  of  LeBlanc  and  Eckardt  in  which  the  persulphate 
is  reduced  by  a  ferrous  salt,  the  process  of  Griitzner,  in  which 
an  arsenite  solution  is  the  reducing  agent,  and  the  arsenate- 
iodide  method,  in  which  the  persulphate  is  determined  by  the 
difference  between  the  amount  of  iodine  in  an  iodide  added  and 
the  amount  necessary  to  oxidize  the  arsenite  remaining  after 
boiling  the  solution,  are  all  in  close  agreement  and  all  higher  than 
those  obtained  by  the  process  of  Namias  or  that  of  Mondolfo. 

The  process  of  LeBlanc  and  Eckardt  is  simple,  rapid  and  con- 
venient. The  method  of  Griitzner  is  advantageous  in  that  the 
ordinary  arsenite  solution  is  the  standard  for  the  process,  though 
requiring  the  application  of  a  correction.  The  arsenate-iodide 
method,  introduced  as  a  control,  is  accurate  but  less  simple 
than  the  other  methods. 

SELENIUM. 

The  Gravimetric  Estimation  of  Selenious  Acid  by  Liberation  of 
Iodine  and  Absorption  of  that  Element  by  Silver. 

When  selenious  acid  is  treated  in  acidulated  solution  with 
potassium  iodide  and  shaken  with  silver  in  an  atmosphere  of 
hydrogen,*  the  reaction  proceeds  to  the  complete  reduction  of 
the  element,  according  to  the  equation 

SeO2  +  4  KI  +  4  HC1  =  4  KC1  +  2  H2O  +  Se  +  2  I2. 

The  increase  in  weight  of  the  insoluble  material  represents  the 
iodine  evolved  plus  the  selenium  of  the  selenium  oxide. f     Tests 

*  See  page  444. 

f  Claude  C.  Perkins,  Am.  Jour.  Sci.,  [4],  xxix,  338. 


376 


METHODS  IN  CHEMICAL  ANALYSIS 


made  in  the  application  of  this  process  to  selenium  dioxide  twice 
crystallized  from  nitric  acid  and  resublimed  over  manganese 
dioxide  are  given  in  the  table. 

Weighing  of  Silver  Iodide  and  Selenium. 


Ag  taken. 

Se  taken. 

Increase. 

Calculated  Se. 

Error. 

grrn. 

grin. 

grm. 

grm. 

grm. 

2.0133 

O  .  0050 

0.0365 

o  .  0049 

—  O.OOOI 

2.0133 
2  -  5639 
2.5639 
3.0018 

0.0075 
0.0126 
o  .  0428 
0.0504 

0.0529 
0.0894 
0.3178 
0.3799 

O  0071 
O.OI2I 
0.0429 
0.0501 

—  O.0004 
—  0.0005 
+  0.0001 

-0.0003 

The  Gravimetric  Determination  of  Selenious  Acid  by  Precipitation 

of  Selenium. 

For  the  gravimetric  determination  of  selenious  acid  it  is  usual 
to  precipitate  the  selenium  with  sulphurous  acid  in  presence  of 
hydrochloric  acid  and  to  weigh  the  elementary  selenium.  Pre- 
cipitation by  this  method,  however,  is  slow  and  incomplete  in 
many  cases,  so  that  it  is  always  necessary  to  treat  the  nitrate  a 
second  time  with  sulphurous  acid  and  to  digest  for  some  time. 
Adopting  the  idea  from  volumetric  methods  for  the  determina- 
tion of  selenium  *  in  which  an  iodide  in  acid  solution  is  used  to 
reduce  the  selenious  acid,  Peirce  f  effects  the  reduction  by  the 
use  of  potassium  iodide  in  large  excess. 

When  potassium  iodide  is  added  in  slight  excess  to  the  solution 
of  selenious  acid  acidified  with  hydrochloric  acid  the  selenium 
is  precipitated  in  the  form  of  a  red  powder.  Boiling  for  ten 
minutes  removes  most  of  the  liberated  iodine  and  changes  the 
selenium  into  the  black  modification.  This  may  be  collected 
upon  an  asbestos  felt,  washed,  dried  at  100°  to  a  constant  weight, 
and  weighed.  If  the  selenium  amounts  to  less  than  o.i  grm. 
the  results  accord  well  with  theory.  When  the  amount  is  larger, 
the  selenium  is  apt  to  assume  on  boiling  a  pasty,  molten  condi- 
tion which  makes  filtering  and  washing  impossible.  In  this 
condition  the  selenium  holds  iodine.  By  using  the  potassium 
iodide  in  large  excess,  however,  the  pasty  condition  may  be 
modified,  and  by  effecting  the  reduction  in  large  volumes  of  solu- 
tion the  danger  of  inclusion  is  lessened. 

*  See  pages  377,  379. 

f  A.  W.  Peirce,  Am.  Jour.  Sci.,  [4],  i,  416. 


SELENIUM 


377 


According  to  this  procedure,  it  is  sufficient  to  dilute  the  solu- 
tion containing  selenious  acid  or  a  selenite  to  400  cm.3  before 
acidifying  with  hydrochloric  acid  and  then  to  add  potassium 
iodide  to  an  amount  about  three  grams  in  excess  of  that  actually 
required.  Boiling  from  ten  to  twenty  minutes  will  change  the 
selenium  to  the  black  modification  and  remove  most  of  the  iodine. 
The  process  of  precipitation  and  filtering  can  be  completed  in 
half  an  hour.  The  selenium  is  dried  at  100°  to  a  constant 
weight. 

When  the  selenium  occurs  in  the  higher  form  of  oxidation  the 
reduction  follows  the  same  course,  though  iodine  is  not  liberated 
until  the  solution  is  quite  warm ;  but  at  the  end  of  the  usual  time 
of  boiling  the  action  is  complete. 

Results  obtained  by  this  procedure  are  given  in  the  table. 

Reduction  of  Selenium. 


Se  taken  as  SeO2. 
grm. 

Se  found, 
grrn. 

KI. 
grm. 

Volume. 
cm.3 

Error, 
grm. 

0.2853 

0.2861 

7 

QOO 

+0.0008 

0.3189 

0.3192 

8 

400 

+0.0003 

0.3318 

0.3324 

7 

500 

+0.0006 

0.3798 

0.3805 

7 

500 

+0.0007 

0.4252 

0.4259 

7 

350 

+0.0007 

0.4430 

0-4434 

10 

450 

+0.0004 

Se  taken  as  SeO3. 

0.1063 

0.1065 

5 

500 

+0.0002 

0.1063 

0.1062 

5 

375 

—  O.OOOI 

O.2OIO 

0.2017 

5 

350 

+0.0007 

0.3H5 

0.3126 

6 

500 

+  0.001  1 

The   lodometric   Determination    of  Selenious   Acid   by   Methods 

Based  upon  the  Action  of  Potassium  Iodide  in  Presence 

of  Acid. 

The  contact  According  to  the  method  of  Muthman  and  Schae- 

Method.  fer^*  wnen  selenious  acid  is  brought  into  contact  with 

potassium  iodide  in  an  acidulated  solution,  iodine  and  selenium 
are  liberated  in  elementary  condition,  the  former  being  directly 
determinable  by  titration  with  sodium  thiosulphate  after  addi- 
tion of  starch.  On  account  of  the  difficulty  in  determining  the 
exact  point  in  the  titration  at  which  the  starch  blue  disappears 
*  Ber.  Dtsch.  chem.  Ges.,  xxvi,  1008. 


378  METHODS  IN  CHEMICAL  ANALYSIS 

from  the  liquid  in  which  the  finely  divided  and  opalescent 
selenium  is  held  in  suspension,  the  process  was  recommended 
by  the  authors  for  use  only  when  great  accuracy  is  not  es- 
sential. 

Evidently  if  the  reaction  between  the  acidulated  iodide  and 
selenious  acid  is  simple  and  complete,  the  process  should  be 
capable  of  improvement  by  removing  the  selenium  before  the 
titration  is  attempted.  This  has  been  done  by  Gooch  and 
Reynolds  *  by  filtration  by  means  of  the  vacuum  pump  upon  a 
thick  felt  of  asbestos  in  a  perforated  crucible  or  cone  of  large 
filtering  surface.  With  a  properly  prepared  filter  of  this  descrip- 
tion there  is  no  difficulty  in  separating  the  selenium  in  a  very  few 
moments  so  completely  that  it  is  possible  to  determine  the 
iodide  remaining  dissolved  in  the  excess  of  potassium  iodide  with 
all  the  accuracy  characteristic  of  this  most  exact  of  titration 
processes.  When,  however,  the  difficulty  of  determining  the 
end-reaction  in  the  titration  of  the  iodine  by  the  thiosulphate  is 
overcome,  it  becomes  apparent  that  the  reaction  itself  is  likely 
to  be  incomplete.  Even  when  the  potassium  iodide  is  used  in 
moderate  excess  in  presence  of  a  large  proportion  of  hydrochloric 
acid,  and  with  addition  of  the  thiosulphate  previous  to  the  filtra- 
tion in  order  that  there  may  be  every  opportunity  for  the  iodine 
and  thiosulphate  to  interact,  the  results  show  marked  deficiency 
in  the  reduction  of  the  selenious  acid.  Either  the  reaction 
according  to  the  equation 

SeO2  +  4  HI  =  Se  +  2  H2O  +  2l2 

is  incomplete  or  else  there  is  formed  between  the  selenium  and 
iodine  some  combination,  such  as  was  noted  by  Hautefeuille  f 
in  the  interaction  between  iodine  and  hydrogen  selenide. 

In  further  study  of  this  reaction,  and  following  the  suggestion 
of  Peirce's  experience  in  the  gravimetric  determination  of  sele- 
nium, Norton  t  has  found  that  the  accuracy  of  the  process  is  very 
much  increased  for  small  amounts  of  selenium  by  the  use  of 
relatively  very  large  amounts  of  potassium  iodide.  The  process 
is  still  inaccurate  when  large  amounts  of  selenium  are  present. 
This  is  shown  by  the  results  of  the  accompanying  table. 

*  F.  A.  Gooch  and  W.  G.  Reynolds,  Am.  Jour.  Sci.,  [3],  1,  255. 

f  Compt.  rend.,  Ixviii,  1554. 

J  J.  T.  Norton,  Jr.,  Am.  Jour.  Sci.,  [4],  vii,  292. 


SELENIUM 


379 


Contact  Method. 


SeO2  used. 

KI. 

Volume  of 
solution. 

HC1 
sp.  gr.  1.  12. 

SeO2  found. 

Error. 

grm. 

grm. 

cm.3 

cm.3 

grm. 

grm. 

0.0553 

IO 

150 

IO 

0.0558 

+0.0005 

0.0574 

5 

ISO 

IO 

0.0567 

—  0.0007 

0.0683 

5 

150 

IO 

0.0683 

o.oooo 

0.0487 

5 

ISO 

10 

o  .  0484 

—0.0003 

0.2617 

10 

150 

10 

0.2589 

—0.0028 

The  Distillation       The  reaction  between  selenious  acid,  potassium 
Method.  iodide,  and  hydrochloric  acid,  in  the  sense  of  the 

equation  given  above,  may  be  pushed  further  toward  completion 
by  submitting  the  mixture  to  distillation.  For  this  purpose 
Gooch  and  Reynolds  *  make  use  of  the  apparatus  previously 
described  and  figured,  f  The  distillation  flask  is  a  Voit  gas- wash- 
ing flask,  and  this  is  sealed  to  the  inlet  tube  of  a  Drexel  wash- 
bottle  used  as  a  receiver,  to  the  outlet  tube  of  which  is  sealed  a 
Will  and  Varrentrapp  absorption  apparatus  to  serve  as  a  trap. 
The  mixture  to  be  distilled,  containing  not  more  than  0.2  grm. 
of  selenium  dioxide,  is  introduced  into  the  flask,  a  solution  of 
3  grm.  of  potassium  iodide  in  100  cm.3  of  water  is  put  into  the 
receiver  and  trap,  and  during  the  distillation  a  slow  current  of 
carbon  dioxide  is  passed  through  the  apparatus  to  keep  the  boil- 
ing regular.  Naturally,  the  acidified  solution  of  the  iodide  in  the 
flask  retains  with  great  tenacity  traces  of  dissolved  iodine,  so 
that,  in  order  to  determine  all  the  iodine  liberated  in  the  reac- 
tion, the  residue  in  the  flask  as  well  as  the  distillate  in  the  receiver 
and  trap  must  be  titrated  with  sodium  thiosulphate.  Results 
are  fairly  good,  though  a  little  deficient,  for  amounts  of  selenium 
dioxide  up  to  0.2  grm.;  but  when  the  amount  of  the  dioxide 
reaches  0.5  grm.  it  is  found  that  the  sum  total  of  iodine  in  the 
^distillate  and  in  solution  in  the  residue  falls  far  below  the  theory 
based  upon  the  assumption  that  the  products  are  selenium, 
iodine,  and  water.  The  details  of  treatment  and  the  results  for 
the  smaller  amounts  of  selenium  afe  recorded  in  the  table. 

The  selenium  in  the  residue  is,  for  the  smaller  amounts,  left, 

after  the  boiling,  in  fine  dense  crystalline  condition,  so  that  it 

does  not  interfere  with  the  titration  of  the  free  iodine;  but  for  the 

larger  amount,  0.5  grm.,  it  is  in  pasty  form  adhering  to  the  flask. 

*  Am.  Jour.  Sci.,  [3],  i,  256.  f  See  Fig.  3,  page  4. 


METHODS  IN  CHEMICAL  ANALYSIS 


Examination  proves  that  the  selenium  holds  iodine,  which  is 
liberated  slowly  to  water  and  more  rapidly  to  an  aqueous  solution 
of  potassium  iodide;  but  the  error  thus  introduced  is  allowable 
up  to  the  limit  of  0.2  grm.  of  selenium  dioxide. 

Distillation  Method. 


SeO2  taken. 

KI  in  flask. 

HClin 
flask 
(sp.  gr. 

1.20). 

Total 
volume 
boiled. 

Time  in 
minutes. 

SeO2  found. 

Error. 

gnu. 

grin. 

cm.3 

cm.1 

grm. 

grm. 

0.0499 

i 

5 

60 

5 

0.0497 

—  O.OOO2 

0.0499 

i 

5 

60 

5 

0.0497 

—  O.OO02 

0.0499 

i 

5 

60 

10 

o  .  0496 

—  O.OO03 

O  .  2OOO 

3 

5 

60 

10 

0.1995 

—  0.0005 

O.2OOO 

3 

5 

60 

10 

«  0.1991 

—  0.0009 

0.2023 

3 

5 

60 

IO 

0.2018 

—  0.0005 

Method:  Treat- 
ment of  the 
Residue. 


Peirce  states  *  that  the  range  of  the  process  may  be  much  ex- 
tended by  the  use  of  very  large  amounts  of  potassium  iodide. 
In  this  case,  however,  more  iodine  is  retained  in  the  residue  and 
the  difficulty  resulting  from  atmospheric  action  when  the  acidu- 
lated residue  is  exposed  becomes  magnified. 
Differential  When  a  solution  of  an  arsenate,  potassium  iodide, 

and  sulphuric  acid  is  boiled  under  defined  conditions! 

arsenic  acid  is  reduced  to  arsenious  acid  with  liber- 
ation of  iodine.  When  the  arsenic  acid  is  in  excess  the  whole  of 
the  iodine  is  evolved  and  the  arsenious  acid  produced  is  its  exact 
measure.  Upon  making  the  solution  alkaline  with  acid  potassium 
carbonate,  the  arsenious  acid  may  be  reoxidized  by  standard 
iodine,  and  the  amount  of  iodine  thus  used  is  the  exact  equivalent 
of  that  set  free  in  the  reduction  process.  Gooch  and  Peirce  J 
have  shown  that  when  selenious  acid  is  present  during  the  reac- 
tion between  arsenic  acid  and  the  iodide,  selenium  is  reduced,  and 
the  subsequent  estimation  of  arsenious  acid  in  the  residue  will 
be  less  than  should  be  produced  by  the  iodide  by  an  amount 
equivalent  to  the  selenious  acid  present,  the  reduction  taking 
place  according  to  the  equation 

SeO2  +  4HI  =  Se  +  2  H2O  +  2  I2. 

*  See  page  ^76. 
t  See  page  457. 
J  F.  A.  Gooch  and  A.  W.  Peirce,  Am.  Jour.  Sci.,  [4],  i,  31. 


SELENIUM 


381 


According  to  the  procedure  worked  out,  the  selenious  acid  to  be 
determined  is  put  into  an  Erlenmeyer  flask  of  300  cm.3  capacity; 
a  known  amount  of  standardized  potassium  iodide  (somewhat 
in  excess  of  that  theoretically  required)  is  added ;  and  a  solution 
containing  about  2  grm.  of  pure  di-hydrogen  potassium  arsenate 
with  20  cm.3  of  sulphuric  acid  of  half-strength  is  introduced. 
During  the  boiling  the  mixture  is  protected  from  ordinary  me- 
chanical loss  by  a  trap  *  (consisting  of  a  two-bulbed  drying  tube, 
cut  short  and  hung  loosely  with  the  wide  end  downward  in  the 
mouth  of  the  flask)  and  by  the  introduction  of  a  few  bits  of  por- 
celain. The  liquid  is  boiled  until  the  volume  decreases,  accord- 
ing to  indicating  marks  on  the  flask,  from  100  cm.3  or  more  to 
35  cm.3,  concentration  to  about  this  lower  limit  having  been 
found  to  be  necessary  for  the  completion  of  the  reaction.  The 
residue  is  cooled,  the  acid  is  nearly  neutralized  with  potassium 
hydroxide,  acid  potassium  carbonate  is  added  until  it  is  present 
to  the  amount  of  20  cm.3  of  its  saturated  solution  in  excess 
of  the  quantity  needed  for  complete  neutralization,  and,  after 
the  addition  of  starch,  standard  iodine  is  introduced  until 
the  starch-blue  appears. 

Differential  Method. 


Initial 
volume. 

Final 
volume. 

H2SO4 
half- 
strength. 

Di-hydro- 
gen- 
potassium 
arsenate. 

KI 

taken. 

800, 

taken. 

Se02 
found. 

sSfSOff 

cm.* 

cm.3 

cm.8 

grm. 

grm. 

grin. 

grm. 

grm. 

IOO 

35 

20 

2 

1.3277 

o.  1280 

0.1275 

—  0.0005 

IOO 

35 

20 

2 

1.0429 

O  .  0998 

0  .  0994 

—  0.0004 

IOO 

35 

20 

2 

1.0887 

O.  1024 

o.  1028 

+0.0004 

IOO 

35 

20 

2 

1.0405 

o.  1036 

0.1028 

—  O.OOO& 

IOO 

35 

2O 

2 

1.0721 

o.  1030 

o.  1029 

—  o.oooi 

IOO 

35 

2O 

2 

0.9958 

0.1273 

o.  1272 

—  O.OOOI 

125 

35 

2O 

2 

2.0828 

0.1997 

O  .  2OOO 

+o  .  0003 

125 

35 

2O 

2 

2.2272 

O.2IIO 

0.2II3 

+0.0003. 

125 

35 

2O 

2 

2.1535 

O.2O67 

0.2069 

+  O.OOO2 

150 

40 

2O 

2 

2-6554 

0.2560 

0.2549 

—  o.ooii 

175 

35 

20 

2 

3.2428 

0.3IIO 

0.3II8 

+0.0008 

175 

35 

2O 

2 

3.2428 

0.3085 

0.3083 

—  0.0002 

The  iodine  introduced  measures  the  arsenious  acid  (and  so 
the  quantity  of  iodine  set  free  by  the  arsenic  acid),  and  the 
difference  between  it  and  the  iodine  originally  present  in  the 

*  See  Fig.  6,  page  6. 


382  METHODS  IN  CHEMICAL  ANALYSIS 

form  of  the  iodide  represents  the  amount  set  free  by  the  seleni- 
ous  acid. 

The  preceding  table  comprises  the  details  and  results  of  a 
series  of  determinations  made  in  the  manner  outlined. 

The  Determination  of  Selenious  Acid  by  Potassium  Permanganate'. 

In  the  action  of  potassium  permanganate  upon  selenious  acid, 
whether  in  a  solution  acidified  with  sulphuric  acid  or  made  alka- 
line by  caustic  soda,  the  reduction  of  the  permanganate  does  not 
proceed  to  the  lowest  degree  of  oxidation  of  the  manganese,  the 
selenious  acid  being  unable  to  reduce  the  higher  hydroxides  which 
separate.  When  the  permanganate  is  introduced  into  the  acid- 
ified solution  the  color  vanishes,  leaving  a  clear  colorless  liquid, 
but  as  more  is  added  the  solution  becomes  yellow  and  deepens 
gradually  in  color  to  a  reddish  brown,  until  turbidity  due  to  the 
deposition  of  a  brown  hydroxide  of  manganese  ensues,  and  finally 
the  characteristic  color  of  the  permanganate  is  plainly  distin- 
guishable. The  exact  point  at  which  precipitation  of  the  man- 
ganic hydroxide  begins  depends  upon  the  dilution,  acidity,  and 
temperature  of  the  solution.  In  employing  the  reaction  quanti- 
tatively it  is  necessary,  according  to  Gooch  and  demons,*  to  add 
the  permanganate  in  distinct  excess,  to  destroy  the  surplus  by 
means  of  standard  oxalic  acid  added  to  the  solution  acidified  with 
sulphuric  acid,  and  then  to  determine  the  excess  of  oxalic  acid 
in  the  warmed  solution  by  addition  of  more  permanganate.  The 
difference  between  the  amount  of  permanganate  actually  used 
and  that  required  to  oxidize  the  known  amount  of  oxalic  acid 
introduced  is  the  measure  of  the  selenious  acid  acted  upon,  pro- 
vided the  amount  of  sulphuric  acid  present  in  the  final  titration 
and  the  temperature  are  adjusted,  to  prevent  on  the  one  hand 
interference  with  the  end-reaction  by  precipitation  of  manganese 
hydroxide,  according  to  Guyard's  reaction,  and  on  the  other 
hand  to  obviate  evolution  of  oxygen  outside  the  main  reaction.! 
According  to  the  method  of  treatment  prescribed,  the  solution  of 
selenium  dioxide  in  100  cm.3  of  water  containing  10  cm.3  of  sul- 
phuric acid  of  half-strength  is  heated  to  75°,  an  approximately 
decinormal  standardized  solution  of  potassium  permanganate  is 
added  until  the  characteristic  color  predominates  over  that  of  the 

*  F.  A.  Gooch  and  C.  F.  demons,  Am.  Jour.  Sci.,  [3],  1,  51. 

t  See  page  47. 


SELENIUM 


383 


brown  hydroxide  deposited  during  the  oxidation,  oxalic  acid  in 
solution  of  known  strength  is  introduced  until  the  excess  of  per- 
manganate has  been  destroyed  and  the  insoluble  hydroxide  dis- 
solved, and,  finally,  at  a  temperature  of  50°  or  less,  permanganate 
is  added  to  coloration.  The  final  volume  varies  between  250  cm.3 
and  350  cm.3,  and  the  sulphuric  acid  (absolute)  between  about 
five  per  cent  at  the  start  and  one  and  a  half  or  two  per  cent 
at  the  end. 

The  determination  of  large  amounts  of  selenious  acid  by  this 
method  is  somewhat  less  advantageous  than  would  be  the  case 
if  the  reduction  of  the  permanganate  proceeded  further  in  the 
first  action.  One  hundred  cubic  centimeters  of  a  standard  solu- 
tion is  as  much  as  can  be  conveniently  handled  in  a  single  process 
of  titration,  and  that  volume  of  decinormal  permanganate  (which 
is  about  as  strong  as  the  standard  solution  should  be  when  accu- 
rate work  is  expected)  is  capable  of  oxidizing  about  0.25  grm.  of 
selenium  dioxide.  In  the  table  are  given  the  results  of  practical 
tests  of  this  method. 

Permanganate  Oxidation. 


Oxygen 

Oxygen 

SeO2  taken. 

equivalent  to 
permanganate 

equivalent  to 
oxalic  acid 

SeO2  found. 

Error. 

used. 

used. 

grm. 

grm. 

grm. 

grm. 

grm. 

O.IOOO 

0.03506 

0.02065 

O.IOOI 

+O  .  OOOI 

O.IOOO 

0.03519 

0.02073 

0.1004 

+O.OOO4 

O.  IOOO 

0.03706 

0.02255 

0.1007 

-f-o  .  0007 

O.IOOO 

0.03853 

0.02422 

0.0994 

—O.OOO6 

O.IOOO 

0.03512 

0.02065 

o.  1005 

+0.0005 

O.  2OOO 

0.06124 

0.03256 

0.1994 

—O.OOO6 

O.  2OII 

o  .  06069 

0.03177 

o  .  2008 

—0.0003 

0.2004 

0.06072 

0.03177 

O.  2OIO 

+0.0006 

O.2O2O 

o  .  06083 

0.03185 

O.2OI2 

—  0.0008 

0.2038 

0.06106 

0.03185 

0.2028 

—  o.ooio 

The  Determination   of  Selenious  Acid   by  the  Direct  Action   of 

Sodium  Thiosulphate,  According  to  the  Method  of  Norris 

and  Fay. 

In  the  method  of  Norris  and  Fay  *  for  the  iodometric  deter- 
mination of  selenious  acid,  advantage  is  taken  of  a  direct  and 
unique  action  of  sodium   thiosulphate   upon   selenium  dioxide 
in  the  presence  of  hydrochloric  acid,  four  molecules  of  sodium 
*  Am.  Chem.  Jour.,  xviii,  703. 


METHODS  IN  CHEMICAL  ANALYSIS 


thiosulphate  acting  upon  one  molecule  of  selenious  acid.*  The 
method,  which  consists  in  treating  the  solution  of  selenious  acid 
in  ice  water,  in  the  presence  of  hydrochloric  acid,  with  an  excess 
of  a  n/io  solution  of  sodium  thiosulphate  and  titrating  back 
the  excess  of  the  thiosulphate  with  iodine,  involves  the  addi- 
tion of  an  excess  of  the  thiosulphate  to  the  solution  of  selenious 
and  hydrochloric  acids,  and  thus  establishes  conditions  which 
demand  care  as  to  the  relation  of  the  acid,  thiosulphate,  degree  of 
dilution,  and  temperature. 

Norton  f  shows  that  with  precautions  noted  the  process  of 
Norris  and  Fay  is  simple,  rapid  and  accurate;  without  ttfem,  as 
the  experimental  results  given  below  indicate,  errors  of  consider- 
able amount  may  enter. 

Reduction  by  Thiosulphate  and  Titration  of  Excess. 


Amount  of  SeO2 
taken. 

gnu. 

HC1 
(sp.  gr.  1.  12). 

cm.* 

Excess  of 
Na2S203. 

cm.s 

SeO2  taken, 
grm. 

Error. 

gnn. 

Volume  at  beginning,  200  cm.3 


0.1042 

5 

24.16 

o.  1041 

—  o.oooi 

0.0611 

10 

13-3 

0.0611 

0.0000 

0.0850 

10 

21.9 

0.0828 

—  0.0022 

0.0757 

25 

13.07 

0.0749 

—0.0008 

0.0540 

25 

21  .02 

0.0522 

—0.0018 

Volume  at  beginning,  400  cm.3 


0.0616 

IO 

2.28 

0.0625 

+0.0009 

0.0628 

10 

7.  II 

0.0631 

+0.0003 

o  .  0508 

10 

11.4 

0.0511 

+0.0003 

0.0587 

10 

12.8 

0.0594 

+0.0007 

0.0807 

IO 

15.3 

0.0813 

+0.0006 

0.0633 

IO 

20.85 

o  .  0638 

+0.0005 

0.0682 

25 

I.  II 

0.0685 

+0.0003 

0.0779 

25 

1-35 

0.0788 

+0.0009 

0.0465 

25 

18.93 

o  .  0469 

+o  .  0004 

It  is  recommended  to  so  adjust  conditions  that  no  more  than 
20  cm.3  of  n/io  thiosulphate  shall  be  present  in  excess.  If  this 
limit  be  placed  upon  the  thiosulphate,  5  cm.3  of  hydrochloric  acid 
(sp.  gr.  1.12)  may  safely  be  present  in  a  volume  of  200  cm.3  at 
the  beginning,  or  10  cm.3  of  the  acid  in  a  volume  of  400  cm.a 
The  presence  of  5  cm.3  of  the  acid  in  400  cm.3  of  solution  is  really 
sufficient  to  bring  about  the  reaction. 

*  The  complete  reaction  is  not  stated. 

t  J.  T.  Norton,  Jr.,  Am.  Jour.  Sci.,  [4],  vii,  287. 


SELENIUM  385 

The  I odometric.  Determination  of  Selenic  Acid  by  the  Action  of  the 

Halogen  Acids. 

Reduction  by  It  has  long  been  known  that  selenic  acid  is  re- 
Add'^hDis-  diicible  by  hydrochloric  acid  with  evolution  of  chlo- 
tiiiation.  rine,  but  the  reaction  was  regarded  as  more  or  less 

uncertain  until  Petterson  showed  *  that  conditions  of  action  may 
be  secured  under  which  the  reduction  proceeds  regularly  accord- 
ing to  the  equation 

SeO3  +  2  HC1  =  SeO2  +  H2O  +  C12. 

The  chlorine  evolved  may  be  estimated  iodometrically  and 
taken  as  the  measure  of  the  selenic  acid  originally  present  or 
of  the  selenious  acid  produced.  According  to  this  method  of 
determination,  it  is  only  necessary  to  boil  a  solution  of  selenic 
acid  in  hydrochloric  acid  of  moderate  concentration,  and  if  the 
solution  is  not  too  dilute  the  reduction  is  obtained  in  a  few 
moments. 

Gooch  and  Evans  |  have  determined  the  limits  within  which  a 
successful  determination  of  the  selenic  acid  may  be  expected. 
It  is  shown  that  so  long  as  the  volume  of  the  hydrochloric  acid, 
sp.  gr.  1. 20,  does  not  amount  to  more  than  10  per  cent  of  the  entire 
liquid  no  chlorine  whatever  is  evolved,  and  that  only  when  the 
percentage  of  this  acid  rises  as  high  as  thirty  does  the  chlorine 
evolved  during  boiling  for  five  minutes  approach  the  theoreti- 
cal yield.  Care  must  be  taken,  however,  not  to  prolong  the  boil- 
ing after  the  solution  reaches  a  concentration  corresponding  to 
hydrochloric  acid  of  half-strength ;  for  under  such  conditions  — 
easily  attained  in  boiling  down  mixtures  of  selenious  acid  and 
hydrochloric  acid  —  over-reduction  may  take  place  and  selenium 
appear  visibly  in  the  distillate.  Obviously  it  is  advantageous,  in 
attempting  the  practical  reduction  of  selenic  acid,  to  begin  the 
distillation  with  acid  of  strength  sufficient  to  insure  the  evolution 
of  chlorine  in  quantity  at  the  outset,  and  it  has  been  found  best 
to  start  with  a  mixture  one- third  of  which  is  the  strongest  aqueous 
hydrochloric  acid,  sp.  gr.  1.20.  With  solutions  so  constituted  the 
reduction  goes  on  rapidly.  Good  results  may  be  expected  when 
the  mixture,  containing  one-third  of  its  volume  of  the  strongest 
aqueous  hydrochloric  acid  at  the  beginning,  is  boiled  until  the 

*  Zeit.  anal.  Chem.  xii,  287. 

t  F.  A.  Gooch  and  P.  S.  Evans,  Jr.,  Am.  Jour.  Sci.,  [3],  1,  400. 


386 


METHODS  IN  CHEMICAL  ANALYSIS 


chlorine  is  expelled,  care  being  taken  that  the  volume  of  the  liquid 
shall  not  become  less  than  two-thirds  of  the  original  volume. 

The  apparatus  *  made  by  sealing  to  the  outlet  tube  of  a  Voit 
wash  bottle  (used  as  a  retort)  to  the  inlet  tube  of  a  Drexel  wash 
bottle  (charged  with  potassium  iodide  and  used  as  a  receiver)  with 
a  set  of  Will  and  Varrentrapp  bulbs  (sealed  to  the  receiver,  to  serve 
as  a  trap)  is  convenient  for  the  operation.  A  current  of  carbon 
dioxide  aids  in  carrying  the  chlorine  to  the  receiver  and  in  pro- 
moting quiet  boiling. 

From  solutions  having  a  total  volume  of  75  cm.3  at  the  outset 
and  containing  25  cm.3  of  the  strongest  aqueous  hydrochloric  acid 
(sp.  gr.  1. 20),  the  entire  amount  of  chlorine  corresponding  to  the 
reduction  of  0.2  grm.  of  selenic  acid  to  selenious  acid  is  liberated 
in  ten  minutes.  The  iodine  in  the  receiver  is  estimated  by  thio- 
sulphate.  The  details  of  experiments  made  under  this  procedure 
with  selenic  acid  obtained  by  oxidizing  with  permanganate  pure 
selenium  dioxidef  are  given  in  the  table. 

Reduction  by  Hydrochloric  Acid. 


SeQa  taken. 

Total  volume 
at  the  outset. 

HCl 
(sp.  gr.  1.20) 
present. 

Time  in 
minutes. 

SeO3  found. 

Error. 

cm.3 

cm.8 

grm. 

0.0572 

75 

25 

IO 

0.0568 

—0.0004 

o  0572 

75 

25 

10 

0.0569 

-0.0003 

o.  1144 

75 

25         . 

10 

0.1143 

—  o.oooi 

o.  1144 

75 

25 

10 

0.1137 

—0.0007 

0.1144 

75 

25 

IO 

0.1147 

+0.0003 

0.2288 

75 

25 

IO 

0.2233 

—0.0005 

0.2288 

75 

25 

IO 

0.2279 

—0.0009 

Reduction  by          When  acted  upon  by  sulphuric  acid  and  potassium 
Hydrobromic      bromide  in  solution,  selenic  acid  liberates  bromine 

Acid,  witn 

Distillation.        in  proportion  to  the  excess  of  acid,  the  amount  of 
bromide,  and  the  temperature. 

SeO3  +  2  HBr  =  SeO2  -f  H2O  +  Br2. 

When  such  a  solution  is  boiled  the  bromine  is  evolved  and  may 
be  collected  in  potassium  iodide,  and  the  iodine  thus  set  free  may 
be  determined  by  standard  sodium  thiosulphate  and  taken  as  the 
measure  of  the  bromine  distilled. 

*  See  Fig.  3,  page  4. 

t  See  page  382. 


SELENIUM 


387 


Gooch  and  Scoville  *  have  shown  that  the  applicability  of  the 
reaction  to  quantitative  purposes  turns  upon  the  adjustment  of 
the  proportions  of  the  reagents  used.  The  apparatus  shown  in 
Fig.  3  t  is  convenient  for  the  distillation  process. 

When  the  proportions  of  sulphuric  acid,  potassium  bromide, 
and  selenic  acid  are  favorable,  the  bromine  liberated  is  removed 
rapidly  to  the  distillate,  leaving  the  residue  perfectly  colorless, 
but  as  the  distillation  is  continued  the  liquid  residue  again  takes 
on  color  and  more  iodine  is  set  free  by  the  action  of  the  distillate 
upon  potassium  iodide,  while  selenium  is  plainly  visible  in  the 
receiver.  When  the  amount  of  potassium  bromide  is  large,  its 
effect  is  to  retain  bromine  in  the  liquid  so  obstinately  that  no 
period  of  colorlessness  intervenes  before  the  second  stage  of  color 
arrives;  when  its  amount  is  small,  while  that  of  the  sulphuric 
acid  is  also  small,  the  reduction  of  the  selenic  acid  and  the  evolu- 
tion of  the  bromine  progress  slowly ;  and  the  interval  of  colorless- 
ness  is  prolonged  when  the  amount  of  bromide  is  small,  while  that 
of  the  acid  is  comparatively  large.  The  proportions  found  best 
in  handling  0.25  grm.  of  selenic  acid,  or  less,  are  an  initial  volume 
of  60  cm.3  containing  20  cm.3  of  sulphuric  acid  of  half-strength, 
with  I  grm.  of  potassium  bromide.  Under  these  conditions  it  is 
found  that  the  reduction  is  almost  theoretically  exact  when  the 
distillation  is  continued  until  the  recoloration  of  the  boiling  liquid 
is  distinctly  recognizable ;  and  this  point  corresponds  in  practice 
very  closely  to  a  concentration  of  volume  to  35  cm.3.  In  the  fol- 
lowing table  are  given  the  results  of  experiments  made  under 
these  conditions  of  action. 

Reduction  by  Hydrobromic  Acid. 


SeO3  taken 
as  H2SeO*. 

HoSO4  of 

half- 
strength. 

KBr 

taken. 

Initial 
volume. 

Final 
volume. 

SeO3 
calculated. 

Error. 

grm. 

cm.8 

grm. 

cm." 

cm.» 

grm. 

grm. 

o  .  0590 

2O 

60 

35 

0.0588 

—  0.0002 

o  .  0590 

20 

60 

35 

0.0591 

-f-o.oooi 

0.0614 

20 

60 

35 

0.0616 

+0.0002 

0.0614 

20 

60 

35 

0.0607 

—  0.0007 

o.  1180 

20 

60 

35 

0.1177 

—0.0003 

0.1180 

20 

60 

35 

o.  ii  80 

o  .  oooo 

0-1534 

20 

60 

35 

0.1527 

—0.0007 

0.2349 

20 

60 

35 

0.2350 

+O.OOOI 

*  F.  A.  Gooch  and  W.  S.  Scoville,  Am.  Jour.  Sci.,  [3],  1,  402. 
f  See  page  4. 


388 


METHODS  IN  CHEMICAL  ANALYSIS 


Reduction  by 


The  determination  of  selenic  acid  by  acting  with 
hydrochloric  acid  and  potassium  iodide  and  estimat- 
with  Distillation.  mg  ^  {od{ne  liberated  has  been  studied  by  Gooch 

and  Reynolds.*  While  the  simple  contact  of  selenic  acid  and 
potassium  iodide  in  solution  acidified  with  hydrochloric  acid  does 
not  produce  a  regular  liberation  of  iodine,  it  is  possible  by  sub- 
mitting such  mixtures  to  distillation,  when  the  amounts  of 
selenic  acid  present  are  not  too  large,  to  bring  about  a  definite 
reaction  in  which  the  products  are  selenium,  water  and  iodine, 
according  to  the  equation 

Se03  +  6HI  =Se  +  3H20  +  3I2. 

In  applying  this  reaction  to  analytical  purposes  it  is  convenient 
to  make  use  of  an  apparatus  and  procedure  previously  described  f 
and  to  treat,  preferably,  not  more  than  0.2  grm.  of  the  selenic 
oxide  with  I  grm.  to  3  grm.  of  potassium  iodide,  5  cm.3  of  con- 
centrated hydrochloric  acid  in  a  total  volume  of  60  cm.3,  and  to 
continue  the  boiling  for  ten  minutes. 

The  results  of  experiments  made  in  this  manner  with  selenic 
acid  obtained  in  solution  by  oxidizing  known  amounts  of  selenium 
dioxide  by  potassium  permanganate  J  are  given  in  the  table. 

Reduction  by  Hydriodic  Acid. 


SeO3  taken. 

Klin 
flask. 

HClin 
flask 
(sp.  gr. 

1.20). 

Total 
volume 
boiled. 

Time  in 
minutes. 

SeO3  found. 

Error. 

grm. 

grm. 

cm.3 

cm  .3 

grm. 

grm. 

0-0593 

i 

5 

60 

5 

0.0593 

0.0000 

0-0593 

i 

5 

60 

5 

0.0591 

—  0.0002 

0-0593 

3 

5 

60 

10 

0.0596 

+o  .  0003 

0.1779 

3 

5 

60 

10 

o.  1769 

—  O.OOIO 

0.1779 

3 

5 

60 

IO 

o.  1780 

+0.0001 

0.1779 

3 

5 

60 

10 

o.  1764 

-0.0015 

Reduction  by  In  a  mixture  made  up  of  a  selenate,  an  arsenate, 
Di^rentLf"*1'  potassium  iodide,  and  sulphuric  acid,  the  arsenic  acid 
Method.  attacks  the  hydriodic  acid  before  all  of  the  selenic  acid 

is  reduced.  In  order  to  apply  to  selenic  acid  the  differential 
method  of  determination,  the  selenic  acid  must  first  be  reduced  to 

*  F.  A.  Gooch  and  W.  G.  Reynolds,  Am.  Jour.  Sci.,  [3],  1,  258. 
t  See  page  379. 
J  See  page  382. 


SELENIUM 


389 


the  condition  of  selenious  acid.  Ordinarily,  the  simplest  mode  of 
reducing  selenic  acid  is  by  boiling  it  in  solution  with  hydrochloric 
acid  of  definite  strength,*  but  in  this  case  the  presence  of  hydro- 
chloric acid  is  precluded  on  account  of  the  consequent  volatiliza- 
tion of  arsenious  chloride  during  the  process  of  concentration  in 
the  subsequent  treatment  with  the  iodide.  It  is  possible,  how- 
ever, to  make  use  of  reduction  by  hydrobromic  acid,  since  arse- 
nious bromide  is  not  appreciably  volatile  under  the  conditions. 
Gooch  and  Peircef  have  shown  that  the  determination  of  selenic 
acid  may  therefore  be  accomplished  by  first  reducing  it  to  selenious 
acid  by  the  bromide  process  and  then  treating  the  residue  by  the 
differential  method  for  the  determination  of  selenious  acid.J 

Differential  Method. 


SeO2  taken  as  H2SeO4. 
grm. 

KI  used  in  second 
reduction. 

grm. 

SeO2  found, 
grm. 

Error, 
grm. 

0.0378 

O  .  6306 

0.0380 

+0.0003 

0.0378 

0.5643 

0.0374 

—0.0004 

0.0516 

0.7136 

0.0517 

+O.OOOI 

o  .  0503 

0.7302 

0.0508 

+o  .  0005 

0.0541 

0.6671 

0.0544 

-fo  .  0003 

o.  1007 

•3277 

O.IOII 

+O.OOO4 

o.  1008 

.3277 

O.IOII 

+0.0003 

0.1007 

.2082 

0.1005 

—  O.OOO2 

o.  1007 

.1684 

0.1016 

+o  .  0009 

o.  1007 

.0522 

0.0999 

—  O.OOOS 

0.1009 

.2679 

o.  1005 

—  0.0004 

0.1031 

1  .1119 

0.1032 

+0.0001 

0.1870 

I  .8720 

0.1879 

+o  .  0009 

0.2014 

1-99*5 

O.2O2O 

+0.0006 

o.  2016 

2.0745 

0.2025 

+0.0009 

0.2059 

1.8687 

o  .  2064 

+o  .  0005 

The  procedure  is  as  follows:  The  solution  containing  the 
selenate  to  be  determined  is  put  in  an  Erlenmeyer  flask  of  300 
cm.3  capacity  with  I  grm.  of  potassium  bromide  and  sulphuric 
acid  amounting  to  20  cm.3  of  the  acid  of  half-strength.  The  so- 
lution amounting  to  60  cm.3  or  100  cm.3  is  boiled  until  the  color- 
less solution  left  when  the  bromine  vanishes  begins  to  color  again. 
Experience  shows  that  the  reappearance  of  the  brownish  color  is 
very  easily  seen  and  that  it  is  not  safe  to  conclude  that  the  free 

*  See  page  385. 

f  F.  A.  Gooch  and  A.  W.  Peirce,  Am.  Jour.  Sci.,  [4],  i,  33. 

J  See  page  380. 


390  METHODS  IN   CHEMICAL  ANALYSIS 

bromine  has  been  eliminated,  under  the  conditions  of  dilution  and 
proportion,  until  this  stage  of  concentration  —  which  corresponds 
to  a  volume  of  about  35  cm.3  —  has  been  reached ;  but  the  dis- 
tillation should  not  be  pushed  beyond  the  point  at  which  the 
returning  color  is  noted.  When  this  condition  has  been  reached 
the  solution  is  cooled  and  treated  exactly  in  the  manner  described 
for  the  reduction  of  selenious  acid.  The  neutralization  by  acid 
potassium  carbonate,  after  the  final  boiling,  generally  occasions 
the  precipitation  of  manganous  carbonate,  but  the  precipitate 
does  not  interfere  in  the  slightest  with  the  titration  which  follows. 
The  preceding  table  comprises  determinations  made  to  test 
the  accuracy  of  the  iodometric  determination  of  selenic  acid  by 
the  combined  processes  of  reduction. 

The  Separation  of  Selenium  from  Tellurium  by  Procedure  Based 
upon  the  Difference  in  Volatility  of  the  Bromides. 

When  small  amounts  of  selenic  acid  are  boiled  in  aqueous  solu- 
tion with  potassium  iodide  and  hydrochloric  acid,  selenium  is 
precipitated,  while  the  iodine  set  free  simultaneously  may  be 
estimated  in  the  distillate  and  residue,  and  taken  as  the  measure 
of  the  selenic  acid  originally  present.*  If  the  iodide  is  omitted 
from  the  mixture,  so  that  the  hydrochloric  acid  alone  shall  be 
the  reducer,  the  reduction  proceeds  only  to  the  point  of  formation 
of  selenious  acid,  provided  the  boiling  is  not  continued  after  the 
hydrochloric  acid  has  reached  the  condition  of  half-strength  at 
which  it  boils  unchanged  under  normal  atmospheric  pressure. 
A  solution  of  selenic  acid,  potassium  bromide,  and  sulphuric 
acid,  of  regulated  dilution  and  proportions,  also  yields  under  de- 
fined conditions  selenious  acid  as  the  product  of  reduction. 
When,  however,  the  ebullition  of  a  solution  of  selenious  acid  in 
hydrochloric  acid  is  continued  after  the  acid  has  reached  the 
condition  of  half -strength,  traces  of  selenium  appear  in  the  receiver 
and  connecting  tubes,  the  distillate  sets  free  iodine  from  potas- 
sium iodide,  and  it  is  evident  that  the  selenious  acid  is  under- 
going further  reduction ;  and  the  same  effects  are  produced  when 
the  boiling  of  the  mixture  of  sulphuric  acid,  potassium  bromide, 
and  selenious  acid  is  pressed  beyond  the  point  at  which  the  solu- 
tion begins  to  be  colored.  Obviously,  under  certain  conditions 

*  See  page  388. 


SELENIUM  391 

of  concentration,  selenium  tetrachloride  and  selenium  tetrabro- 
mide,  respectively,  are  forming  from  the  acid;  and  the  appear- 
ance of  the  elementary  selenium  is  due  to  partial  decomposition 
of  the  halogen  salts.  Phenomena  of  a  similar  nature  are  seen 
when  an  aqueous  solution  of  selenious  acid,  phosphoric  acid,  and 
sodium  chloride  is  submitted  to  distillation :  that  is  to  say,  there 
comes  a  time  in  the  process  of  boiling  such  mixtures  when  the 
appearance  of  elementary  selenium  and  the  action  of  the  distil- 
late upon  potassium  iodide  make  evident  the  volatilization  and 
partial  decomposition  of  the  selenium  compounds  of  the  halo- 
gens, and  the  further  continuance  of  the  treatment  results  in  the 
more  or  less  complete  removal  of  the  selenium  compounds  to  the 
distillate.  From  the  mixture  containing  the  phosphoric  acid, 
selenious  acid  and  sodium  chloride  only  a  partial  volatilization 
of  the  selenium  chloride  takes  place.  The  volatilization  of 
selenium  bromide,  however,  produced  by  the  reaction  between 
phosphoric  acid,  selenious  acid  and  potassium  bromide,  may  be 
made  complete.  Upon  this  reaction  Gooch  and  Peirce  *  have 
based  a  process  for  the  separation  of  selenium  and  tellurium, 
taking  advantage  of  the  volatility  of  selenium  tetrabromide 
and  the  non-volatility  of  tellurium  tetrabromide  under  definite 
conditions. 

The  distillation  apparatus  used  in  the  process  of  separation 
is  shown  in  Fig.  4-f  It  consists  of  two  Voit  flasks,  a  Drexel 
bottle,  and  Will  and  Varrentrapp  bulbs,  connected  by  sealed  or 
ground  joints,  as  shown  in  the  figure. 

The  operation  is  conducted  as  follows:  In  the  first  Voit  flask, 
V1,  selenium  dioxide  and  tellurium  dioxide  are  dissolved  in  potas- 
sium hydroxide,  the  alkali  is  neutralized  and  the  precipitate  thus 
formed  is  redissolved  by  phosphoric  acid,  added  in  excess  to  the 
amount  of  20  cm.3  of  the  acid  of  sp.  gr.  1.70.  To  the  solution 
is  added  i  grm.  of  potassium  bromide  with  enough  water  to 
make  the  entire  volume  of  the  solution  50  cm.3  The  second  flask, 
V2,  contains  10  cm.3  of  water,  and  the  Drexel  bottle  and  trap  are 
charged  with  a  solution  of  potassium  iodide.  Carbon  dioxide  is 
passed  through  the  apparatus  and  the  solution  in  V1  is  boiled  until 
the  volume  has  diminished  to  15  cm.3,  the  flask  and  connecting 
tube  being  cloaked  with  a  mantle  of  asbestos  board  and  gently 

*  F.  A.  Gooch  and  A.  W.  Peirce,  Am.  Jour.  Sci.,  [4],  i,  181. 
t  See  page  5. 


392 


METHODS  IN  CHEMICAL  ANALYSIS 


flamed  toward  the  last  to  remove  traces  of  the  selenium  bromides 
held  back  mechanically  by  the  oily  tellurium  compound  which 
collects.  After  cooling,  the  first  flask  V1  is  removed;  I  grm.  of 
potassium  iodide  and  5  cm.3  of  hydrochloric  acid  are  added  to 
the  contents  of  the  second  flask,  V2;  the  current  of  carbon  dioxide 
is  again  started  through  the  apparatus;  the  mixture  is  boiled 
ten  minutes;  and  the  iodine  in  the  flask,  receiver  and  trap, 
determined  by  titration  with  sodium  thiosulphate,  is  taken  as 
the  measure  of  the  selenium  dioxide.* 

Results  of  this  procedure  are  given  in  the  table. 

Double  Distillation. 


Te02 

taken. 

grm. 

KBr 

taken. 

grm. 

H3P04 

(sp.  gr.  1.70) 
taken. 

cm.3 

Final 
volume. 

cm.8 

SeO2  taken, 
grm. 

SeO2  found. 

grm. 

Error, 
grin. 

2O    . 

15 

o  .  0366 

0.0372 

+o  .  0006 

20 

15 

0.0366 

0.0377 

+O.OOII 

2O 

15 

o.  1098 

o.  1090 

—  O.OOOS 

20 

15 

o.  1098 

O.IIOI 

+  O.OOO3 

O.I 

20 

15 

0.0733 

0-0735 

+  O.OOO2 

0. 

2O 

15 

0.0997 

0.0995 

—  O.OOO2 

O. 

2O 

15 

o.  1004 

o.  1003 

—  O.OOOI 

O. 

2O 

15 

0.0916 

0.0914 

—  0.0002 

0. 

20 

15 

0.0997 

0.0995 

—  O.OOO2 

0. 

20 

15 

O.  IOIO 

o.  1014 

-J-O.OOO4- 

O. 

2O 

15 

o.  1015 

0.1008 

—  O.OOO7 

O. 

2O 

IS 

o.  1019 

O.  IO22 

+  0.0003 

O. 

2O 

IS 

O.  IOIO 

O.  IOI2 

+  O.OOO2 

0. 

2O 

15 

O.  IOO2 

O.  IOOO 

—  O.OO02 

0. 

20 

15 

0.1006 

o.  1004 

—  O.OOO2 

O. 

2O 

15 

o.  1006 

O.IOOI 

—  0.0005 

The  phenomena  of  the  distillation  are  very  characteristic. 
When  selenious  acid  is  present  without  tellurous  acid,  the  solu- 
tion boils  quietly  in  the  first  flask  until  the  volume  of  liquid 
has  decreased  to  about  30  cm.3,  when  traces  of  red  selenium  begin 
to  deposit  in  the  tube  joining  the  first  and  second  flask.  When 
the  volume  has  further  diminished  to  about  25  cm.3  the  liquid 
begins  to  take  on  color,  darkens  rapidly,  and  evolves  bromine, 
which  at  once  attacks  the  selenium  previously  deposited.  The 
greater  part  of  the  bromine  is  absorbed  in  the  second  flask,  V2, 
but  a  trace  finds  its  way  to  the  Drexel  bottle,  in  which  it  sets  free 
a  slight  amount  of  iodine  from  the  iodide.  As  the  operation 

*  See  page  379. 


SELENIUM  393 

progresses,  an  orange-yellow  crystalline  solid,  presumably  sele- 
nium tetrabromide  for  the  most  part,  appears  in  the  tube  where 
the  selenium  has  been,  while  a  dark  oily  liquid,  consisting 
largely,  no  doubt,  of  the  monobromide,  condenses  in  drops  upon 
the  walls  of  the  flask  and  returns  to  form  a  floating  layer  upon 
the  hot  liquid.  Finally,  when  the  volume  has  diminished  to 
15  cm.3,  the  liquid  has  become  perfectly  clear  and  colorless,  white 
fumes  of  hydrobromic  acid  are  evolved,  and  the  tube  between  the 
two  flasks  has  been  cleared.  At  this  point  the  second  flask,  V2, 
contains  (besides  a  trace  of  selenium  corresponding  to  the  slight 
amount  of  bromine  which  has  escaped  to  the  Drexel  bottle)  the 
colorless  selenious  acid  regenerated  by  the  action  of  the  water 
and  free  bromine  upon  the  mixed  selenium  bromides.  The 
contents  of  this  flask  may  now  be  treated  with  potassium  iodide 
and  hydrochloric  acid  as  directed  above  *  and  the  iodine  in  the 
receiver,  including,  of  course,  the  small  amount  set  free  by  the 
bromine  which  reaches  the  receiver  in  the  first  stage  of  the  process, 
and  the  small  amount  remaining  in  the  flask,  measure  the  selenium 
dioxide  acted  upon. 

When  tellurium  dioxide  is  subjected  without  the  selenium 
dioxide  to  similar  treatment  the  phenomena  are  different.  The 
solution  containing  the  tellurous  acid,  potassium  bromide,  and 
phosphoric  acid,  in  the  proportions  used  in  the  experiments  with 
selenious  acid,  colors  at  about  the  same  degree  of  concentration 
at  which  the  solution  containing  the  selenious  acid  began  to 
darken.  As  the  concentration  progresses,  the  color  deepens, 
ruby  red  crystals  (probably  hydrated  tellurium  tetrabromide) 
form,  which  accumulate  upon  the  walls  of  the  flask  and  turn 
yellow,  and  when  the  volume  of  the  solution  is  diminished  to 
15  cm.3  a  green  vapor  begins  to  distil.  During  the  process  no 
iodine  is  set  free  in  the  Drexel  bottle,  and  upon  stopping  the 
boiling  and  adding  potassium  iodide  to  V2  no  iodine  is  liberated, 
even  when  the  boiling  has  gone  so  far  that  a  trace  of  the  green 
vapor  has  condensed  and  run  into  the  water  in  the  flask. 

When  the  tellurium  dioxide  and  selenium  dioxide  are  both 
present  the  characteristic  phenomena  occur  together  —  the  evo- 
lution of  bromine,  coloring  of  the  liquid,  distillation  of  selenium 
bromides,  crystallization  of  tellurium  tetrabromide,  and  volatiliza- 
tion of  the  selenium  compounds. 

*  See  page  379. 


394 


METHODS  IN  CHEMICAL  ANALYSIS 


TELLURIUM. 

The  Gravimetric  Estimation  of  Tellurous  Acid  by  Liberation  of 
Iodine  and  Absorption  of  that  Element  by  Silver. 

Tellurous  acid  reacts  with  potassium  iodide  in  presence  of 
hydrochloric  acid  and  silver  *  in  an  atmosphere  of  hydrogen 
according  to  the  reaction 

TeO2  +  4  KI  +  4  HC1  =  4  KC1  +  2  H2O  +  Te  +  2  I2. 

The  increase  in  weight  of  insoluble  material  represents  both  iodine 
and  tellurium.  Tests  made  by  Perkins  |  with  tellurium  dioxide 
prepared  from  the  basic  nitrate  gave  the  following  results. 

Weighing  of  Silver  Iodide  and  Tellurium. 


Ag  taken, 
grni. 

Te  taken, 
grm. 

Increase, 
grm. 

Calculated  Te. 
grm. 

Error, 
grm. 

2.0152 

0.0330 

0.1654 

0.0332 

+O.OOO2 

2.0152 

O  .  0990 

Q-4931 

o  .  0989 

—  0.0001 

2.0815 

0.0528 

0.2635 

0.0529 

+O.OOOI 

2.0815 

o  .  0660 

0.3294 

0.0661 

+  O.OOOI 

2.0815 

O  .  0990 

0.4948 

o  .  0993 

+o  .  0003 

2.1693 

0.1650 

0.8240 

0.1654 

+o  .  0004 

3.0126 

o.  1650 

0.8258 

0.1657 

+0.0007 

3.0126 

o  .  0660 

0.3302 

o  .  0663 

+0.0003 

The  Determination  of  Tellurous  Acid  by  Oxidation  with  Potassium 

Permanganate. 

The  estimation  of  tellurous  acid  by  oxidation  with  excess  of 
potassium  permanganate  (either  in  acid  or  alkaline  solution), 
destruction  of  the  higher  oxides  of  manganese  or  the  manganate 
by  standard  oxalic  acid  in  presence  of  sulphuric  acid,  and  titra- 
tion  of  the  residual  oxalic  acid  by  more  permanganate  has  been 
shown  by  Brauner  J  to  be  feasible,  but  the  tendency  of  the 
'permanganate  to  throw  off  too  much  oxygen  when  the  oxidation 
is  made  in  solutions  strongly  acidified  with  sulphuric  acid  (as 
must  be  the  case  if  the  tellurous  oxide  is  to  be  held  perma- 
nently in  solution  by  sulphuric  acid)  necessitates  the  application 
of  a  considerable  correction.  Gooch  and  Danner  §  have  shown, 

*  See  page  444. 

f  Claude  C.  Perkins,  Am.  Jour.  Sci.,  [4],  xxix,  540. 

t  Jour.  Chem.  Soc.,  1891,  238. 

§  F.  A.  Gooch  and  E.  W.  Danner,  Am.  Jour.  Sci.,  [3],  xliv,  301. 


TELLURIUM 


395 


however,  that  when  the  tellurous  oxide  is  first  dissolved  in  an 
alkali  hydroxide  and  the  solution  is  made  acid  to  a  limited  degree 
with  sulphuric  acid,  either  before  or  after  oxidation  by  the 
permanganate,  no  correction  appears  to  be  necessary. 

According  to  the  first  procedure,  the  alkaline  solution  of  the 
oxide  is  diluted  to  100  cm.3,  a  measured  amount  of  standardized 
permanganate  is  added  in  excess,  sulphuric  acid  [i  :  i]  is  intro- 
duced to  an  amount  not  exceeding  by  more  than  5  cm.3  that 
needed  for  neutralization,  standardized  oxalic  acid  is  measured 
in  to  an  amount  more  than  sufficient  to  destroy  the  manganic 
oxide  and  permanganate  left  after  the  oxidation,  and  the  surplus 
of  oxalic  acid  is  titrated  by  permanganate. 

According  to  the  second  procedure,  the  alkaline  solution  of  the 
oxide  is  treated  with  sulphuric  acid  [i  :  i]  until  the  precipitate 
first  thrown  down  is  just  redissolved,  and  I  cm.3  more  of  the 
[i  :  i]  acid  is  added.  To  the  solution  thus  acidulated  is  measured 
in  potassium  permanganate  in  excess,  and  then  standardized 
oxalic  acid  to  an  amount  a  little  more  than  sufficient  to  destroy 
the  permanganate  remaining;  the  liquid  is  warmed  to  80°,  and 
the  excess  of  oxalic  acid  is  titrated  by  permanganate. 

The  results  of  experiments  made  according  to  these  procedures 
are  given  below. 

Oxidation  in  Alkaline  Solution. 


TeO2  taken, 
grin. 

TeO2  found, 
grm. 

Error, 
grm. 

Mean  error, 
grm. 

0.1200 

0.1199 

—  O.OOOI        1 

0.0783 

0.0783 

0  .  0000 

0.0931 
O.  IIOO 

o  .  0938 
o.  1116 

+0.0007      I 
+0.0016       1 

l.         +0.0006 

o  .  0904 

o  .  0907 

+0.0003 

0.1065 

0.1077 

+O.OOI2        J 

Oxidation  in  Acid  Solution. 


TeO2  taken, 
grm. 

TeO2  found, 
grm. 

Error, 
grm. 

Mean  error, 
grm. 

0.0910 

0.0912 

+O.OOO2       "1 

0.0910 

O  .  0908 

—  O.OOO2 

0.0911 
0.0913 

0.0922 
0.0913 

+O.OOII          I 
O.OOOO         | 

+0.0003 

0.0912 

0.0913 

+0.0001 

0.0914 

0.0921 

+0.0007     J 

396  METHODS  IN  CHEMICAL  ANALYSIS 

Oxidation  in  ^n  tne  presence  of  free  hydrochloric  acid  the  action 

Presence  of  a  of  the  permanganate  upon  tellurous  acid  has  been 
shown  by  Brauner  *  to  be  irregular  and  excessive, 
and  the  irregularity  cannot  be  corrected  (as  in  the  titration  of 
ferrous  salts  in  presence  of  hydrochloric  acid)  by  the  addition  of 
a  manganous  salt  according  to  the  well-known  procedure  of 
Kessler  |  and  Zimmermann.t  Gooch  and  Peters  §  have  pointed 
out  that  there  should  be  nothing  to  prevent  the  accurate  deter- 
mination of  tellurium  in  tellurous  compounds  in  the  presence  of 
chlorides  by  the  permanganate  process,  provided  the  first  oxi- 
dation is  made  in  alkaline  solution  and  the  second  oxidation  is 
carried  out  with  such  precautions  as  are  necessary  to  a  correct 
determination  of  oxalic  acid  by  permanganate  in  presence  of 
hydrochloric  acid;  for  the  special  danger  of  over-action  on  the 
part  of  the  permanganate  cannot  exist  while  the  solution  is 
alkaline,  and  has  passed  when  the  tellurite  has  become  a  tellurate 
and  before  the  solution  is  made  acid. 

It  has  been  shown  that  the  presence  of  a  manganous  salt  is 
necessary  and  sufficient  to  secure  regularity  of  action  when 
oxalic  acid  is  titrated  in  presence  of  a  considerable  amount  of 
hydrochloric  acid.  When  the  amount  is  no  more  than  would 
be  formed  in  the  decomposition  of  a  gram  or  two  of  halogen  salt 
of  tellurium  the  disturbing  effect  under  ordinary  conditions  of 
work  is  probably  inappreciable,  but  even  in  such  a  case  it  is 
better  to  oxidize  in  the  presence  of  a  manganous  salt  for  the 
reason  that  the  titration  of  the  oxalic  acid  may  then  be  made 
at  the  ordinary  atmospheric  temperature. 

According  to  the  procedure  recommended,  the  alkali  hydroxide 
solution  of  tellurous  oxide  containing  the  alkali  chloride  is  treated 
with  standardized  potassium  permanganate  until  the  character- 
istic permanganate  color  is  visible;  standardized  ammonium 
oxalate  is  introduced  in  excess  of  the  quantity  required  to  reduce 
the  remaining  permanganate,  manganate,  and  higher  oxides;  and 
enough  sulphuric  acid  [i  :  i]  is  added  to  neutralize  the  alkali 
hydroxide  and  leave  an  excess  of  about  5  cm.3.  Then  the  final 
titration  with  permanganate  may  be  made  either  after  heating 

*  Jour.  Chem.  Soc.,  1891,  241. 

t  Ann.  Phys.,  cxviii,  48;   cxix,  225,  226. 

I  Ann.  Chem.,  ccxiii,  302. 

§  F.  A.  Gooch  and  C.  A.  Peters,  Am.  Jour.  Sci.,  [4],  viii,  122. 


TELLURIUM 


397 


the  solution  to  80°  or  at  the  ordinary  temperature  after  the  ad 
dition  of  0.5  grm.  to  I  grm.  of  manganous  chloride. 

Experimental  results  of  the  procedure  are  given  below: 

Permanganate  Oxidation  in  Alkaline  Solution:    Treatment  with  Oxalic  Acid 
Permanganate  Titration  in  Acid   Solution. 

Volume  at  beginning,  150  cm.3:  Te  =  i27.5. 


TeO2  taken, 
grm. 

NaCl. 

grm. 

H2SO4  1  :  i. 
cm.J 

MnCl2.4H20. 
grm. 

TeO2  found, 
grm. 

Error, 
grm. 

Temperature  of  titration,  6o°-8o°  C. 


O.IOOO 

0.4 

5 

0.1003 

+0.0003 

O.IOOO 

0.4 

5 

O.IOOO 

0.0000 

O.IOOO 

0.4 

5 

0.1004 

+0.0004 

O.IOOO 

i  .0 

5 

.  .  . 

0.1003 

+0.0003 

0.0650 

i  .0 

5 

0.0653 

+0.0003 

Temperature  of  titration,  2o°-26°  C. 


0.0700 

0.4 

5-7 

I.O 

0.0705 

+0.0005 

0.0700 

0.4 

5-7 

I.O 

0.0698 

—  0.0002 

0.0700 

0.4 

5-7 

0-5 

0.0701 

+0  .  0001 

O.IOOO 

0.4 

5-7 

0-5 

o.  1008 

+0.0008 

Oxidation  in  Fairly  good  determinations  of  tellurous  acid  may 

Presence  of  a      be  made  similarly  in  the  presence  of  a  bromide,  pro- 
vided   the    titration   is   made   at   the    atmospheric 
temperature  in  the  presence  of  a  sufficiency  (0.5  grm.  to  I  grm.) 
of  a  manganous  salt  and  of  an  excess  of  sulphuric  acid  limited  to 

Permanganate  Oxidation  in  Alkaline  Solution:    Treatment  with  Oxalic  Acid: 
Permanganate  Titration  in  Acid  Solution. 

Volume  at  beginning,  150  cm.3:   Te  =  i27.5. 
Temperature  of  titration,  24°-26°  C. 


TeO2  taken. 

NaCl. 

KBr. 

H2S04 
12.5  per 
cent. 

MnCl2. 
4H20. 

TeO2  found. 

Error. 

grm. 

grm. 

grm. 

cm.8 

grm. 

grm. 

grm. 

0.0650 

o-5 

I 

I.O 

0.0661 

+O.OOII 

0.0650 

0-5 

I 

I  .O 

0.0647 

—  0.0003 

O.IOOO 

o-5 

I 

I  .O 

O.  IOO2 

+O.OOO2 

0.3000 

o-5 

5 

o-5 

O.3OIO 

+0.0010 

0.0650 

0-5 

o-5 

i 

i  .0 

0.0661 

+O.OOII 

398  METHODS  IN  CHEMICAL  ANALYSIS 

5  cm.3  of  the  12.5  per  cent  mixture,  in  a  volume  of  150  cm3. 
At  higher  temperatures  and  higher  concentrations  of  acid,  bro- 
mine is  liberated  by  the  permanganate.  The  experimental  results 
are  given  in  the  table. 

The  Determination   of   Tellurous    Acid  by  the  Precipitation    of 
Tellurous  Iodide. 

Hydriodic  acid  and  tellurous  acid  interact  with  the  formation 
of  tellurium  tetraiodide,  converted  by  water  into  an  oxyiodide 
and  by  excess  of  alkali  iodides  to  soluble  double  salts.  Gooch 
and  Morgan  *  have  observed  that  when  potassium  iodide  is  added 
to  a  cold  solution  of  tellurous  acid  containing  at  least  one-fourth 
of  its  volume  of  strong  sulphuric  acid,  no  tendency  to  form  a 
double  salt  becomes  apparent  until  the  potassium  iodide  amounts 
to  more  than  enough  to  convert  all  the  tellurous  acid  present  into 
tellurium  tetraiodide  according  to  the  equation 

H2Te03  +  4  H2S04  +  4  KI  =  TeI4  +  4  KHSO4  +  3  H2O. 

The  tellurium  tetraiodide  thus  formed  is  extremely  insoluble  in 
sulphuric  acid  of  one-fourth  strength,  though  soluble  in  excess  of 
potassium  iodide  and  acted  upon  by  water  with  formation  of 
tellurium  oxyiodide  and  hydriodic  acid.  It  is  produced  at  first 
in  the  condition  of  a  finely  divided  dark  brown  precipitate  which 
upon  agitation  of  the  liquid  gathers  in  curdy  masses  and  settles, 
leaving  the  liquid  clear.  By  taking  advantage  of  this  tendency 
to  curd,  it  is  possible  to  determine  without  great  difficulty  the 
exact  point  during  the  gradual  addition  of  potassium  iodide  when 
the  precipitation  of  the  tellurium  iodide  is  complete.  Upon  this 
property  Gooch  and  Morgan  have  based  a  simple  titrimetric 
method  for  the  direct  determination  of  small  amounts  of  tellu- 
rous acid. 

According  to  the  procedure  described,  tellurous  oxide  is  dis- 
solved in  a  very  little  of  a  strong  solution  of  potassium  hydroxide, 
and  dilute  sulphuric  acid  is  added  carefully  until  the  tellurous 
acid  which  is  precipitated  upon  neutralization  of  the  alkali 
hydroxide  is  just  redissolved.  To  this  solution,  contained  in  an 
Erlenmeyer  flask,  sulphuric  acid  [i  :  i]  is  added  in  such  amount 
that  the  liquid  shall  contain,  after  the  subsequent  addition  of 
potassium  iodide  in  solution,  at  least  one-fourth  of  its  volume  of 
*  F.  A.  Gooch  and  W.  C.  Morgan,  Am.  Jour.  Sci.,  [4],  ii,  271. 


TELLURIUM 


399 


strong  sulphuric  acid.  The  flask  is  placed  upon  a  pane  of  window 
glass  supported  upon  strips  of  wood  about  I  cm.  above  the  level 
of  a  work  table  covered  with  white  paper.  A  solution  of  approx- 
imately decinormal  potassium  iodide,  free  from  iodate  and  care- 
fully standardized  in  terms  of  iodine  by  a  method  to  be  described,* 
is  introduced  gradually  from  a  burette  into  the  middle  of  the 
Erlenmeyer  beaker.  As  the  drops  of  the  potassium  iodide  touch 
the  liquid  the  precipitation  forms  at  the  center  and  travels  in 
rings  toward  the  outer  walls  of  the  beaker.  When  the  liquid  be- 
comes so  opaque  that  the  effect  of  the  potassium  iodide  is  dis- 
tinguished with  difficulty,  the  beaker  is  rotated  and  the  curded 
precipitate  permitted  to  settle ;  and  then  the  process  of  titration 
is  continued  as  before  until  precipitation  ceases.  With  an  Erlen- 
meyer flask  10  cm.  in  diameter  across  the  bottom  and  a  final 
volume  of  liquid  amounting  to  not  more  than  100  cm.3  the  pre- 
cipitation is  easily  followed. 

The  results  of  a  series  of  determinations  made  according  to  the 
method  described  are  recorded  in  the  following  table : 

Precipitation  of  Tellurous  Iodide. 
Te  =  i27* 


Final 
volume. 

cm.3 

Strongest 
H2SO4  present. 

cm.3 

Iodine  value 
of  KI  used. 

grm. 

TeOj  taken, 
grm. 

TeO2  found, 
gnu. 

Error, 
grm. 

50 

I? 

o  .  0706 

0.0223 

O.O22I 

—  O.OOO2 

5° 

17 

o  .  0764 

0.0244 

0.0239 

—  0.0005 

50 

17 

O.ISQI 

o  .  0496 

0.0499 

+0.0003 

60 

17 

0.1655 

0.0517 

0.0519 

+O.OOO2 

60 

17 

0.1578 

o  .  0498 

0.0494 

—  O.OOO4 

80 

3° 

0.1591 

o  .  0498 

0.0499 

+0.0001 

100 

30 

0.3179 

O.  IOOI 

0.0997 

—  0.0004 

100 

30 

0.3186 

o.  1008 

0.0999 

—  0.0009 

IOO 

30 

0.3208 

O.  IOII 

0.1005 

—  0.0006 

IOO 

30 

0.3208 

O.  IOIO 

0.1005 

-0.0005 

Determined  by  permanganate  oxidations  and  reductions  by  hydrobromic  acid  (see  p.  402). 


The  lodometric  Estimation  of  Tellurous  Acid. 

The  determination  of  tellurous  acid  by  oxidation  of  the  alkali 
hydroxide  solution  with  potassium  permanganate,  reduction  of 
residual  permanganate  and  higher  oxides  of  manganese  with 
oxalic  acid  in  presence  of  sulphuric  acid,  and  titration  of  the  excess 

*  See  page  457. 


400  METHODS  IN  CHEMICAL  ANALYSIS 

of  oxalic  acid  by  permanganate,  is  not  feasible  in  presence  of  an 
iodide,  because  upon  acidifying  the  mixture,  iodine  is  at  once 
set  free. 

Potassium  permanganate  and  the  higher  oxides  of  manganese 
are,  however,  completely  and  rapidly  reduced  by  an  excess  of 
potassium  iodide  upon  the  addition  of  acid,  and  the  iodine  liber- 
ated is  a  measure  of  the  permanganate.  Norris  and  Fay  *  have 
utilized  this  reaction  in  an  excellent  iodometric  method  for  the  de- 
termination of  tellurous  acid.  This  method  consists  in  treating 
the  alkaline  solution  of  tellurous  oxide  with  standard  permanga- 
nate until  the  meniscus  of  the  liquid  shows  a  deep  pink  color, 
then  diluting  the  solution  with  ice  water,  adding  potassium  iodide 
and  sulphuric  acid,  and  titrating  with  sodium  thiosulphate.  The 
difference  between  the  amount  of  iodine  thus  found  and  the 
amount  found  by  treating  similarly  the  same  amount  of  perman- 
ganate, taken  by  itself,  is  the  measure  of  the  tellurous  acid. 

It  is  plain  that  any  agent  capable  of  converting  the  iodine  to 
hydriodic  acid  without  at  the  same  time  reducing  telluric  acid 
should  be  capable  of  measuring  the  excess  of  the  permanganate, 
and  so  the  amount  of  tellurous  acid  originally  present.  Gooch 
and  Peters  f  make  use  of  the  standard  arsenite,  employed 
also  in  standardizing  the  permanganate,!  to  take  up  the  free 
iodine. 

According  to  this  procedure,  the  solution  of  tellurous  oxide  in 
alkali  hydroxide  is  added  to  a  solution  of  potassium  iodide; 
standardized  permanganate  is  run  in  until  the  green  color  of  the 
manganate  appears  (about  30  cm.3  of  the  n/io  solution  for  every 
o.i  grm.  of  TeC>2);  dilute  sulphuric  acid  is  introduced  in  slight 
excess,  followed,  after  the  iodine  has  separated,  by  an  excess  of 
acid  potassium  carbonate ;  and  the  iodine  is  titrated  to  vanishing 
color  (without  starch)  by  standard  arsenite.  It  is  evident  that 
when  the  solution  is  acidified  more  than  enough  iodide  to  com- 
plete the  reduction  of  the  manganese  oxides  should  be  present,  or 
else  that  the  arsenious  acid  should  be  present  in  suitable  amount 
before  the  sulphuric  acid  is  put  in.  This  latter,  procedure  may 
be  used  in  case,  for  any  reason,  it  is  preferred  not  to  introduce 
more  iodide  into  the  solution  than  was  present  originally,  as, 

*  Am.  Chem.  Jour.,  xx,  278. 

f  F.  A.  Gooch  and  C.  A.  Peters,  Am.  Jour.  Sci.,  [4],  viii,  125. 

j  See  page  41. 


TELLURIUM 


401 


for  example,  when  a  direct  determination  of  the  iodine  present  is 
to  follow. 

Experimental  results  follow  in  the  table: 

Permanganate  Oxidation  in  Alkaline  Solution  and  lodometric  Determination  of 

the  Excess. 


TeOj 
taken. 

NaCl. 

KBr. 

KI. 

Total 
volume 
at  end. 

NaOH 

present 
during 
oxidation. 

TeO2 
found. 

Error. 

gnn. 

grm. 

grm. 

gnu. 

cm.8 

grm. 

grm. 

grm. 

0.1000 

o-S 

1  60 

O.I 

0.1005 

+0.0005 

0.1000 

o-5 

1  60 

O.I 

O.  IOOI 

+0.0001 

O.IOOO 

. 

o-S 

160 

O.I 

0.1003 

+0.0003 

O.  IOOO 

i  .0 

250 

O.I 

o.  1007 

+0.0007 

O.2OOO 

I.O 

250 

0.2 

0.1997 

-0.0003 

O.  IOOO 

0.5 

0-5 

o-S 

250 

O.I 

O.  IOOO 

o.oooo 

0.2100 

1.0 

1.0 

I.O 

225 

0.2 

0.2105 

+0.0005 

O.IOOO 

o-S 

1  60 

I.O 

O.  IOII 

+O.OOII 

0.2000 

I.O 

300 

2.O 

o  .  2009 

+0.0009 

The  lodometric  Determination  of  Telluric  Acid. 

Gooch  and  Rowland  *  have  shown  that  telluric  acid  may 
be  reduced  by  the  action  of  potassium  bromide  and  sulphuric 
acid  to  the  condition  of  tellurous  acid  and  estimated  by  deter- 
mining the  iodine  liberated  by  the  bromine  set  free  in  the 
operation. 

According  to  the  method  demonstrated,  the  alkali  tellurate  is 
introduced  into  the  apparatus  for  distillation  with  3  grm.  of 
potassium  bromide,  care  being  taken  to  insure  in  the  50  cm.3 
or  more  of  liquid  the  presence  of  10  cm.3  of  sulphuric  acid  of 
half  strength.  A  current  of  carbon  dioxide  is  passed  through 
the  apparatus,  and  the  solution  is  boiled  to  set  free  the  bromine, 
which  is  absorbed  in  potassium  iodide  and  estimated  by  standard 
sodium  thiosulphate. 

The  distillation  apparatus  f  consists  of  a  Voit  gas-washing 
flask  which  is  joined  by  a  sealed  joint  to  the  inlet  tube  of  a 
Drexel  washing  bottle.  To  the  outlet  tube  of  the  Drexel  bottle 
is  sealed  a  Will  and  Varrentrapp  absorption  apparatus.  The 

*  F.  A.  Gooch  and  J.  Rowland,  Am.  Jour.  Sci.,  [3],  xlviii,  375. 
t  See  Fig.  3,  page  4. 


4O2 


METHODS  IN  CHEMICAL  ANALYSIS 


washing  bottle  and  attached  bulbs  contain  a  solution  of  3  grm. 
of  potassium  iodide,  and  the  former  is  kept  cool  by  standing  it 
during  the  distillation  in  a  vessel  of  cold  water. 

The  formation  of  tellurium  tetrabromide  in  the  concentrated 
acid  liquid  makes  it  impossible  to  tell  by  the  color  when  all  the 
bromine  has  been  distilled,  but  the  evidence  of  the  experiments 
goes  to  show  that  the  boiling  of  the  liquid  from  a  volume  of 
50  cm.3  to  25  cm.3  is  sufficient,  while  concentration  from  100  cm.3 
to  20  cm.3  apparently  does  no  harm. 

The  tellurite  used  in  testing  this  procedure  was  made  from 
tellurous  oxide  which  had  been  shown  by  titration  with  perman- 
ganate to  have  an  equivalent  weight  of  about  159,  which  corre- 
sponds to  an  atomic  weight  of  127  for  the  element  tellurium. 
Results  are  given  in  the  table. 

Reduction  by  Hydrobromic  Acid  and  Estimation  of  Bromine  Set  Free. 


Initial  volume. 
cm.» 

Final  volume, 
cm.* 

TeO2  taken, 
grm. 

TeO2  found, 
grm. 

Error 
grm. 

5° 

2O 

O.OIO2 

0.0098 

—  O.OOO4 

50 

20 

O.OIO2 

0.0099 

—  0.0003 

5° 

2O 

O.OI02 

o  .  0098 

—  0.0004 

5° 

20 

O.OI02 

0.0098 

—  0.0004 

100 

40 

O.  IOOO 

0.0994 

—O.OOO6 

80 

40 

O.IOOI 

O.  1001 

o.oooo 

IOO 

2O 

O.IOO2 

0.  IOOI 

—  O.OOOI 

50 

2O 

O.  IOOO 

0.1003 

+o  .  0003 

50 

25 

0.50II 

0.5008 

-0.0003 

50 

25 

0.5002 

o  .  5006 

+o  .  0004 

50 

25 

O.5OOO 

0.4998 

—  O.OOO2 

50 

2O 

0.5000 

0.4994 

—  0.0006 

The  Precipitation  of  Tellurium  Dioxide  and  the  Separation  of 
Tellurium  from  Selenium. 

Browning  and  Flint  *  have  shown  that  fairly  accurate  and  con- 
cordant estimations  of  tellurium  may  be  obtained  by  acting  upon 
the  alkaline  solution  of  a  tellurite  with  hydrochloric  acid,  am- 
monium hydroxide,  and  acetic  acid,  sucessively,  and  weighing  the 
precipitate  as  tellurium  dioxide,  the  best  of  all  the  forms  in  which 

*  Philip  E.  Browning  and  William  R.  Flint,  Am.  Jour.  Sci.,  [4],  xxviii, 

112. 


TELLURIUM  403 

tellurium  has  been  weighed.  It  is  unaffected  by  the  air,  is  anhy- 
drous, is  not  hydroscopic,  and  can  easily  be  obtained  in  pure 
condition.  It  can  be  heated  to  any  temperature  under  that  of 
low  redness  without  any  danger  of  volatilization. 

All  processes  for  the  estimation  of  tellurium  in  which  the 
tellurium  is  precipitated  and  weighed  in  elementary  condition 
are  open  to  the  objections,  first,  that  there  is  more  or  less  difficulty 
in  securing  completeness  of  precipitation  owing  to  the  rapid  in- 
crease of  free  acid  *  in  the  solution;  and,  second,  that  the  product 
is  extremely  susceptible  to  oxidation.  On  the  other  hand,  those 
methods  in  which  compounds  decomposable  by  heat  are  trans- 
formed to  the  dioxide  by  ignition  are  generally  both  tedious  by 
reason  of  the  length  of  time  required  (as,  for  example,  the  basic 
nitrate  process  as  described  by  Norris  f)  and,  what  is  more  to 
the  point,  liable  to  errors  caused  not  only  by  lack  of  constancy 
of  composition  but  also  by  the  volatilization  of  the  product  to 
be  weighed.  The  process  of  Browning  and  Flint  presents  there- 
fore special  advantages  in  the  determination  of  tellurium. 

According  to  the  preferred  procedure  set  forth,  the  material 
is  dissolved  in  hydrochloric  acid  or  in  a  10  per  cent  solution 
of  potassium  hydroxide,  about  2  cm.3  for  0.2  grm.  of  dioxide. 
From  the  solution  acidified  with  hydrochloric  acid  and  diluted 
with  boiling  water  to  a  volume  of  200  cm.3  the  finely  crystalline 
tellurium  dioxide  is  precipitated  by  the  careful  addition  of  dilute 
ammonia  in  faint  excess  followed  by  the  faintest  possible  excess 
of  acetic  acid. 

If  these  simple  operations  are  properly  carried  out,  the  pre- 
cipitate will  become  crystalline  by  the  time  the  alkali  hydroxide 
is  in  excess ;  the  addition  of  a  few  drops  of  acetic  acid  causes  the 
precipitation  to  become  entirely  quantitative  when  the  solution 
has  cooled,  so  that  no  tellurium  will  be  found  in  the  filtrate  by 
stannous  chloride.  The  precipitate  can  be  transferred,  and  safely 
and  rapidly  washed  with  cold  water,  and  dried  to  constant 
weight  at  about  105°  (or  even  under  low  redness)  in  a  quarter 
of  an  hour.  Furthermore,  the  filtration  can  be  performed  at 
the  end  of  half  an  hour,  or  after  twenty-four  hours,  as  most 
convenient. 

*  Crane,  Am.  Chem.  Jour.,  xxiii,  409.  See  also  Lenher  and  Homburger, 
Jour.  Am.  Chem.  Soc.,  xxx,  387. 

t  Jour.  Am.  Chem.  Soc.,  xxviii,  1675. 


404  METHODS  IN  CHEMICAL  ANALYSIS 

Test  results  are  given  below : 

The  Solution  in  HCl  Diluted  with  Boiling  Water,  and  Treated  with  Ammonia 
and  Acetic  Acid. 


TeO2  taken, 
grni. 

TeO2  found, 
gnu. 

Error, 
grm. 

0.2002 

0.2000 

—  0.0002 

0.20IQ 

O.2OI7 

—  O.OOO2 

0.2904 
O.2OO6 

O.2OO2 
O.2OO4 

—  O.OOO2 
—  O.OOO2 

O.2OTI 

O.  2OIO 

—  O.OOOI 

0.2003 

0.2003 

0.0000 

The  Solution  in  KOH  Acidified  with  HCl,  Diluted  with  Boiling  Water, 
Treated  with  Ammonia  and  Acetic  Acid. 


2  TeO2.HNO3 
taken.* 

TeO2  theory: 
Te  taken  as  127.5. 

TeO2  found. 

Error. 

grm. 

grm. 

grm. 

grm. 

0.2502 

o  .  2089 

0.2083 

—  O.OOo6 

0.2524 

0.2108 

O.  2IIO 

+0.0002 

0.2505 

o.  2092 

o  .  2089 

—0.0003 

0.2528 

O.  2III 

0.2106 

—0.0005 

0.2531 

0.2113 

0.2106 

—0.0007 

0.5008 

0.4182 

0.4182 

0.0000 

0.5010 

0.4183 

0.4175 

—0.0008 

0.5005 

0.4179 

0.4178 

—  O.OOOI 

Dissolved  in  KOH. 


Separation  from  Selenium. 

If  hydrochloric  acid  solutions  of  tellurium  and  selenium  diox- 
ides be  mixed,  abundantly  diluted  with  boiling  hot  water,  and 
the  operation  of  the  above  described  process  properly  applied, 
only  the  tellurium  is  precipitated,  the  selenium  remaining  en- 
tirely in  solution  in  the  filtrate.  This  not  only  provides  a  simple 
and  rapid  preparative  process  for  the  purification  of  tellurium 
from  selenium,  but  also  makes  possible  the  estimation  of  tellu- 
rium directly  in  the  presence  of  the  latter  element.  The  dilution 
should  be  made,  however,  with  boiling  hot  water,  as  cold  water 
induces  a  flocky  precipitation  and  inclusion  of  selenious  acid. 
Results  are  given  in  the  table. 


TELLURIUM 


405 


Separation  of  Tellurium  from  Selenium 


TeO2  taken. 

grm. 

SeO2  taken, 
grm. 

TeO2  found, 
grm. 

Error, 
grm. 

The  solution  faintly  acid  with  HC1,  diluted  cold. 

0.2000 

o.  2015 
0.2038 

O.I 
O.I 
O    I 

0  .  2002 

o.  2016 
o  .  2040 

+O.OOO2 
+O.OOOI 
+0   0002 

The  solution  faintly  acid  with  HC1,  diluted  hot  and  treated  with  NH4OH 

and  HOC2H3O. 


0.2028 

0.05 

O.2OIQ 

—  O.OOO9 

0.2024 

0.05 

o.  2024 

O   OOOO 

o  .  2003 

O.I 

0.1992 

—  o.oon 

o  .  2009 

O.I 

0.2003 

—0.0006 

CHAPTER  X. 
CHROMIUM;  MOLYBDENUM;  URANIUM. 

CHROMIUM. 

The  Estimation  of  Chromium  as  Silver  Chromate. 

It  has  been  shown  by  Autenrieth  *  that  when  chromic  acid  is 
added  to  a  boiling  solution  of  silver  nitrate,  or  when  a  soluble 
chromate  or  dichromate  is  added  to  a  solution  of  silver  nitrate 
previously  acidified  with  nitric  acid,  or  when  silver  chromate  is 
treated  with  nitric  acid,  silver  dichromate  is  formed;  and  that, 
on  the  other  hand,  it  is  silver  chromate  which  is  precipitated 
when  silver  nitrate  in  excess  is  added  to  a  solution  of  a  soluble 
dichromate,  cold  or  hot,  the  reaction  proceeding  according  to 
the  equation 

4  AgN03  +  K2Cr207  +  H2O  =  2  Ag2CrO4  +  2  KNO3  +  2  HNO3. 

The  characteristics  of  both  silver  dichromate  and  silver  chro- 
mate have  recently  been  summarized  and  further  studied  by 
Margosches.t  The  solubility  of  silver  dichromate  in  water  and 
in  ordinary  solutions  is  such  as  to  preclude  the  use  of  this  sub- 
stance as  the  final  product  of  a  quantitative  process  depending 
upon  precipitation.  The  solubility  of  silver  chromate  in  a 
moderately  large  volume  of  water  is  also  considerable,  and  the 
solvent  action  of  free  acid,  even  acetic  acid  in  quantity,  is  marked. 
Gooch  and  Weed  %  have  found,  however,  that  the  precipitation 
of  silver  chromate  is  practically  complete  in  a  solution  only 
faintly  acid  with  acetic  acid  and  in  presence  of  a  large  excess  of 
silver  nitrate.  If  such  a  precipitate  is  collected  in  the  filtering 
crucible  and  washed  with  a  dilute  solution  of  silver  nitrate  until 
no  other  impurities  remain,  silver  chromate  does  not  dissolve,  and 
the  excess  of  silver  nitrate  may  be  removed  by  the  cautious  use 
of  water  without  appreciable  effect  upon  the  precipitate. 

*  Ber.  Dtsch.  chem.  Ges.,  xxxv,  2057. 
t  Zeit.  anorg.  Chem.,  xli,  68;  1,  231. 

t  F.  A.  Gooch  and  L.  H.  Weed,  Am.  Jour.  Sci.,  [4],  xxvi,  85. 
.406 


CHROMIUM 


407 


From  the  results  of  test  experiments  it  is  apparent  that  accu- 
rate determinations  of  chromium  taken  as  the  chromate  or 
dichromate  may  be  secured  by  precipitating  silver  chromate  in 
presence  of  an  excess  of  silver  nitrate,  making  the  solution  am- 
moniacal  and  then  faintly  acid  with  acetic  acid,  transferring  the 
precipitate,  after  standing  half  an  hour,  to  the  filtering  crucible, 
washing  with  a  dilute  solution  of  silver  nitrate,  and,  after  other 
soluble  impurities  have  been  removed,  finishing  the  washing  with 
small  amounts  of  water  applied  in  successive  portions. 

The  results  are  given  in  the  table.  In  no  case  did  the  filtrate, 
with  the  washings,  show  by  the  lead  acetate  test  the  presence 
of  a  chromate. 

Precipitation  of  Silver  Chromate. 


Ag2CrO4. 

K2Cr207 

AgNO3  used 

Volume  at 

taken. 

in  precipitation. 

precipitation. 

Found. 

Theory. 

Error. 

gnn. 

grm. 

cm.3 

grm. 

grm. 

grm. 

0.0921 

0.4248 

TOO 

o.  2072 

O.  2076 

—0.0004 

0.0921 

0.4248 

IOO 

0.2073 

O.  2076 

—  0.0003 

0.0921 

0.4248 

IOO 

0.2075 

o.  2076 

—  o.oooi 

0.0921 

O  .  4248 

IOO 

o.  2074 

o.  2076 

—  O.OOO2 

0.0921 

0.4248 

100 

0.2075 

o.  2076 

—o.oooi 

0.0921 

0  .  4248 

IOO 

0.2073 

o.  2076 

—0.0003 

0.0921 

0.4248 

ICO 

0.2073 

o.  2076 

—0.0003* 

o  0921 

0.4248 

IOO 

0.2075 

o.  2076 

—  O.OOOI* 

0.0921 

0.4248 

IOO 

O.2O8O 

o.  2076 

-fo  .  0004! 

0.0921 

0.4248 

TOO 

0.2070 

0.2076 

—  o.ooo6f 

0.5801 

3 

150 

1.3087 

1.3082 

+0.0005 

0-7352 

3 

2OO 

1.6573 

1-6574 

—  o.oooif 

*  The  precipitation  was  made  in  presence  of  5  grm.  of  N  H^NOj. 
t  The  precipitation  was  made  in  presence  of  5  grm.  of  XaNO2. 
t  An  excess  of  I  cm.3  of  40  per  cent  acetic  acid  was  added  before  filtering. 

The  lodometric  Determination  of  Chromic  Acid. 

Kessler  *  has  shown  that  arsenious  acid  may  be  estimated  by 
treating  it,  in  the  presence  of  hydrochloric  acid,  with  an  excess 
of  a  chromate  solution  of  known  strength,  by  which  treatment 
the  arsenious  acid  is  oxidized  and  the  chromic  acid  reduced,  and 
determining  the  excess  of  chromic  acid  by  adding  an  excess  of 
a  ferrous  salt  and  titrating  with  chromic  acid  until  a  drop  taken 
from  the  solution  fails  to  give  a  blue  color  with  a  ferricyanide. 

*  Ann.  Phys.,  xcv,  204. 


408  METHODS  IN  CHEMICAL  ANALYSIS 

The  amount  of  the  chromate  originally  used  less  the  excess  de- 
termined by  the  ferrous  salt  gives  the  amount  of  the  chromate 
used  for  the  oxidation,  from  which  may  be  calculated  the  amount 
of  arsenious  acid  originally  present.  Despite  the  use  of  a  ferrous 
salt  and  the  numerous  steps  involved  in  the  manipulation,  Kessler 
claims  very  satisfactory  results  for  his  method.  Browning  *  gives 
results  of  experiments  showing  that  Kessler's  reaction  may  be 
used  in  the  reverse  process  for  the  determination  of  chromic  acid, 
the  arsenious  acid  being  used  in  excess,  according  to  the  reaction 

4  CrO3  +  3  As2O3  +  (*)As2O3  =  2.  Cr2O3  +  3  As2O5  +  (*) As2O3. 

According  to  the  procedure  indicated,  an  excess  of  n/io 
arsenite  is  added  to  the  cold  solution  of  chromic  acid  acidulated 
with  10  cm.3  of  dilute  hydrochloric  acid  or  sulphuric  acid  [i  13], 
the  total  volume  being  less  than  100  cm.3  After  a  few  minutes, 
when  the  solution  has  taken  on  the  bluish  green  color  character- 
istic of  chromic  salts,  the  solution  is  treated  with  acid  potassium 
carbonate  or  acid  sodium  carbonate  in  excess  (about  5  grm.). 
To  the  alkaline  solution  n/io  iodine  is  added  in  excess,  the  mix- 
ture is  allowed  to  stand,  with  frequent  shaking,  for  about  half 
an  hour,  residual  iodine  is  taken  up  with  n/io  arsenite,  and  the 
excess  of  the  last  is  titrated  by  n/io  iodine  in  presence  of  starch. 

The  long  period  of  standing  is  made  necessary  by  the  tendency 
of  precipitated  chromic  hydroxide  to  hold  arsenious  acid  and  thus 
delay  the  oxidation  by  iodine.  The  precipitation  of  the  chromic 
hydroxide  may  be  obviated  by  addition  of  Rochelle  salt  before 
the  neutralization,  but  in  this  event  the  dark  green  color  taken 
on  by  the  solution  makes  the  end-reaction  of  the  starch  iodide 
difficult  to  determine  with  great  accuracy. 

In  the  presence  of  a  ferric  salt  the  brown  color  of  the  precipi- 
tate also  makes  the  determination  of  the  end-reaction  difficult 
unless  the  precipitate  is  allowed  to  settle  after  each  addition  of 
iodine.  Edgar  f  has  shown,  however,  that  this  difficulty  may  be 
obviated  by  adding  sirupy  phosphoric  acid  (3  cm.3  to  5  cm.3) 
before  the  neutralization,  so  that  the  iron  precipitate  is  white 
and  the  starch  blue  conies  out  against  the  pale  green  of  the  pre- 
cipitate. 

Results  of  test  determinations  are  given  in  the  table. 

*  Philip  E.  Browning,  Am.  Jour.  Sci.,  [4],  i,  35. 
t  See  page  511. 


CHROMIUM 


409 


Reduction  by  Standard  Ar senile. 


CrO3  taken, 
grin. 

CrO3  found, 
grm. 

Error, 
grm. 

Remarks. 

O.  IOOI 

o.  1004 

+O  .  0003 

0.1005 

0.1004 

—  O.OOOI 

The  iodine  acted  20  minutes. 

o.  1006 

o.  1007 

+O.OOOI 

The  iodine  acted  20  minutes. 

O.IOO| 

O.  IOII 

+0.0007 

The  iodine  acted  20  minutes. 

o.  1009 

O.  IOOQ 

o  .  oooo 

The  iodine  acted  2  hours. 

O.  IO02 

0.1003 

+O.OOOI 

The  iodine  acted  2  hours. 

O.  IOII 

o.  1004 

—0.0007 

Rochelle-  salt  used. 

o.  1007 

o.  1007 

o  .  oooo 

Rochelle  salt  used. 

0.0401 

0-0395 

—0.0006 

o  .  0402 

o  .  0388 

—0.0014 

0.5  grm.  ferric  alum  present. 

O.  IOOI 

o.  1018 

+o.  1017 

o.  1009 

o.  1007 

—  O.OOO2 

o.  1007 

O.  IOII 

+0.0004 

i  grm.  ferric  alum  present. 

0.1005 

o.  1017 

+O.OOI2 

0.1004 

O.  IOIO 

+o  .  0006 

O.  IOOO 

0.1032 

+0.0032 

Rochelle  salt  used. 

0.1005 

o.  1006 

+O.OOOI 

i  grm.  ferric  alum  present. 

The  lodometric  Estimation  of  Chromic  Acid  and  Vanadic  Acid. 

That  vanadic  acid  and  chromic  acid  may  be  accurately  esti- 
mated in  presence  of  one  another  by  taking  advantage  of  the 
differential  reducing  actions  of  hydrobromic  and  hydriodic  acids 
has  been  shown  by  Edgar.* 

In  carrying  out  the  operation,  the  alkali  salts  of  the  chromic 
and  vanadic  acid  are  put  in  the  Voit  flask  of  the  distillation 
apparatus  previously  described,!  one  or  two  grams  of  potassium 
bromide  are  added,  the  flask  is  connected  with  the  absorption 
apparatus  containing  a  solution  of  potassium  iodide  made  alka- 
line with  sodium  carbonate  or  sodium  hydroxide,  and  the  whole 
apparatus  is  filled  with  hydrogen  gas.  Fifteen  to  twenty  cubic 
centimeters  of  concentrated  hydrochloric  acid  are  added  through 
the  separatory  funnel  and  the  solution  is  boiled  for  ten  minutes, 
an  interval  of  time  found  to  be  enough  for  the  completion  of 
the  reduction.  A  slow  current  of  hydrogen  is  maintained  to 
avoid  back  suction  of  the  liquid  from  the  Drexel  bottle.  The 
apparatus  is  disconnected,  the  Voit  flask  placed  in  a  beaker 
containing  cold  water,  and  the  alkaline  solution  in  the  absorp- 
tion apparatus  cooled  by  running  water.  The  contents  of  the 

*  Graham  Edgar,  Am.  Jour.  Sci.,  [4],  xxvi,  333. 
t  See  Fig.  3,  page  4. 


4io 


METHODS  IN   CHEMICAL  ANALYSIS 


trap  are  washed  into  the  Drexel  bottle  and  the  solution  therein 
is  made  slightly  acid  with  hydrochloric  acid.  The  liberated  iodine 
is  titrated  with  approximately  n/io  sodium  thiosulphate  and  the 
color  is  brought  back  by  a  drop  or  two  of  n/io  iodine  solution, 
after  the  addition  of  starch. 

Alkaline  potassium  iodide  is  again  placed  in  the  absorption 
apparatus  and  the  latter  connected  with  the  Voit  flask.  The 
current  of  hydrogen  is  turned  on  and,  after  the  air  has  been  ex- 
pelled, the  apparatus  is  disconnected  momentarily,  one  or  two 
grams  of  potassium  iodide  are  added  to  the  solution  in  the  Voit 
flask,  and  connections  made  again.  Through  the  separatory  fun- 
nel 10  cm.3  to  15  cm.3  of  concentrated  hydrochloric  acid  and  3  cm.3 
of  sirupy  phosphoric  acid  are  added  and  the  solution  in  the 
reduction  flask  is  boiled  to  a  volume  of  10  cm.3  to  12  cm.3.  The 
absorption  apparatus  is  removed  and  cooled,  hydrochloric  acid  is 
added  and  the  liberated  iodine  titrated  with  approximately  n/io 
sodium  thiosulphate. 

Double  Treatment  with  Hydrobromic  Acid  and  with  Hydriodic  Acid. 


V2O6  taken 
NaV03. 

CrO3  taken 
as 
K2Cr207. 

I. 

Titration. 
Na2S203 
w/ioXi.031. 

II. 
Titration. 
Na2S203 
w/ioXi.o3i. 

Error  on  V2O5. 

Error  on 
CrO3. 

Rrm. 

grm. 

cm.3 

cm.3 

grm. 

grm. 

O    1^23 

16.  20 

16.  22 

(  I)  (      o.oooo 

*»«  ^-o^O 

(II)   I   +0.0002 

O    152? 

16.  IQ 

16.  20 

(  I)  (  —  o.oooi 

w  •     0    O 

•*•    •  xy 

(II)   (        0.0000 

(  I)  j  —  o.oooi 

o  .  203  i 

21  .  <Q 

21  .  IQ 

oy 

oy 

(II)   (    —0.0002 

0.1523 

0.0685 

36.08 

l6.22 

+0.0002 

—  O.OOOI 

0.1523 

o  0685 

36.10 

16.  20 

o  .  oooo 

o.oooo 

0.1523 

o  0085 

36.12 

16.  17 

—0.0003 

+O.OOO2 

0.1523 

o  0685 

36.07 

16.  20 

o  .  oooo 

—  O.OOOI 

0.1523 

0.1370 

56.00 

16.17 

—0.0003 

+O.OOOI 

0.1523 

0.1370 

56.02 

l6.  22 

+O.OOO2 

o.oooo 

0.1523 

0.1370 

56-03 

16.19 

—  O.OOOI 

+0.0001 

W/IOX0.992,  W/IOX0.992 


0.2031 

0.1370 

63.82 

22.46 

0  .  OOOO 

—  O.OOOI 

0.2031 

0.1370 

63.80 

22.48 

+O.OOO2 

—0.0003 

0.1016 

0.0685 

32.00 

11.25 

+O.OOO2 

+O.OOOI 

0.1016 

0.0685 

31.92 

11.24 

+O.OOOI 

—o.oooi 

0.1016 

0.0685 

31.90 

11.25 

+O.OOO2 

—  O.OOO2 

0.0508 

0.0343 

15-95 

5-63 

+O.OO02 

—o.oooi 

0.0508 

0.0343 

15-95 

5-62 

-j-o.oooi 

—  O.OOOI 

CHROMIUM 

The  iodine  determined  in  the  first  titration  corresponds  to  a 
reduction  of  the  chromic  and  vanadic  acids  according  to  the 
equation 

V2O5  +  2  CrO3  +  8  HBr  =  V2O4  +  Cr2O3  +  4  Br2  +  4  H2O, 

while  in  the  second  case  the  iodine  corresponds  to  a  reduction  of 
the  vanadium  tetroxide  to  trioxide  as  indicated  in  the  equation 

V2O4  +  2  HI  =  V2O3  +  I2  +  H2O. 

The  second  titration,  therefore,  determines  the  vanadic  acid 
present,  and  the  difference  between  the  first  and  second  furnishes 
the  necessary  data  for  the  calculation  of  the  chromium. 

Results  obtained  in  the  application  of  this  method  to  the  deter- 
mination of  the  acidic  oxides  contained  in  sodium  vanadate  and 
potassium  dichromate  are  given  in  the  table. 

The  Estimatioh  of  Chromic  Acid  and  Vanadic  Acid  by  Reductions 

and  Oxidations. 

For  the  estimation  of  both  vanadic  acid  and  chromic  acid,  when 
present  together,  the  following  method  has  been  worked  out  by 
Palmer:* 

Into  a  measured  portion  of  the  solution,  made  acid  with  hy- 
drochloric acid,  sulphur  dioxide  is  passed  until  the  reduction 
of  the  vanadium  and  the  chromium  is  complete;  the  solution 
is  then  boiled  in  a  current  of  carbon  dioxide  until  the  last  traces 
of  sulphur  dioxide  are  expelled.  To  the  cooled  solution  a  suffi- 
cient excess  of  potassium  ferricyanide  and  potassium  hydrox- 
ide are  added  in  solution.  By  this  process  the  chromium  is 
oxidized  from  the  condition  of  Cr2O3  to  the  condition  of  CrO3  and 
the  vanadium  from  the  condition  of  V2O4  to  the  condition  of 
V2O5.  After  allowing  the  solution  to  stand  a  few  minutes,  a 
solution  of  barium  hydroxide  is  added  to  complete  precipitation. 
The  combined  precipitates  of  barium  chromate  and  barium  vana- 
date are  filtered  off  on  asbestos  and  thoroughly  washed ;  the 
filtrate  is  made  acid  with  dilute  hydrochloric  acid  and  titrated 
with  a  known  amount  of  permanganate  in  excess,  and  the  excess 
titrated  with  n/2O  potassium  ferrocyanide  to  permanent  green 
coloration  in  presence  of  a  trace  of  ferric  salt. 

*  Howard  E.  Palmer,  Am.  Jour.  Sci.,  [4],  xxx,  141. 


412 


METHODS  IN  CHEMICAL  ANALYSIS 


In  another  portion  of  the  solution  the  vanadium  is  deter- 
mined as  follows:  The  solution,  about  100  cm.3  in  volume,  is 
made  acid  with  from  10  cm.3  to  15  cm.3  of  glacial  acetic  acid,  and 
hydrogen  peroxide  is  added.  The  solution  is  then  heated  to 
boiling  and  boiled  for  a  few  minutes;  by  this  process  the  per- 
chromic  acid  and  the  pervanadic  acid,  which  were  first  formed 
in  the  cold,  are  decomposed,  the  chromium  being  reduced  to  the 
condition  of  Cr2O3,  while  the  vanadium  appears  in  the  condi- 
tion of  V2O5.  The  solution  is  diluted  somewhat,  and  a  solution 
of  lead  acetate  is  added  to  complete  precipitation  of  the  lead 
vanadate;  the  chromium,  being  in  the  condition  of  Cr2O3,  is  not 
precipitated.  The  solution  is  stirred  vigorously  and  heated  to 
boiling  to  coagulate  the  precipitate.  The  precipitate  is  filtered 
off  on  asbestos,  washed  thoroughly  and  dissolved  in  potassium 
hydroxide,  and  the  solution  in  potassium  hydroxide  is  made 
strongly  acid  with  sulphuric  acid,  whereby  the  lead  is  precipitated 
as  the  sulphate,  while  the  vanadic  acid  remains  in  solution.  A 
current  of  sulphur  dioxide  is  passed  through  the  solution  until 
the  blue  color  indicates  complete  reduction  of  the  vanadium  to 
the  tetroxide  condition;  and  the  sulphur  dioxide  is  expelled  by 
boiling  in  a  current  of  carbon  dioxide.  The  warm  solution  is 
then  titrated  with  permanganate  to  the  appearance  of  the  first 
permanent  pink  color,  easily  recognized  in  the  presence  of  the 
white  precipitate  of  lead  sulphate. 

Determination   of  Chromic  Acid  and  Vanadic  Acid. 


V,Oj  taken, 
grin. 

CrO3  taken, 
grm. 

V2O6  found, 
grm. 

Error, 
grm. 

CrOj  found, 
grm. 

Error, 
grm. 

0.1139 

O.IOIO 

O.II34 

—0.0005 

O.  IOIO 

O.OOOO 

0.1139 

O.IOIO 

O.II39 

o.oooo 

o.  1017 

+o  .  0007 

0.1139 

O.  IOIO 

O.H34 

—  o.oooo 

o.  1019 

+o  .  0009 

0.1139 

O.IOIO 

O.II42 

+o  .  0003 

o.  1019 

+o  .  0009 

0.1139 

O.  IOIO 

O.II3I 

—0.0008 

o.  1015 

+o  .  0005 

0.1139 

O.  IOIO 

O.II34 

—0.0005 

o.  1016 

+o  .  0006 

0.1139 

0.0505 

O.II39 

O  .  OOOO 

0.0507 

+  O.OOO2 

0.1139 

0.0505 

O.II34 

—0.0005 

0.0507 

+  O.OOO2 

0.0569 

0.0505 

0.0565 

—0.0004 

0.0505 

O  .  OOOO 

0.0569 

0.0505 

0.0563 

—0.0006 

0.0508 

+0.0003 

This  titration  gives  a  measure  of  the  amount  of  vanadium 
present;  and  by  subtracting  the  number  of  cubic  centimeters  of 
permanganate  used  in  this  titration  from  the  number  of  cubic 


CHROMIUM  413 

centimeters  used  in  the  preceding  titration,  the  number  of  cubic 
centimeters  corresponding  to  the  oxidation  of  the  chromium  from 
Cr2O3~to  CrOs  is  obtained.  Experimental  results  are  given  in 
the  table. 

The  Volumetric  Estimation  of  Chromium  in  the  Chromic  Condition  ~ 

As  Bollenbach  and  Luchmann  have  shown,*  chromium  may  be 
quantitatively  oxidized  from  the  condition  of  Cr2O3  to  the  condi- 
tion of  CrO3  by  potassium  ferricyanide  in  alkaline  solution,  and 
a  measure  of  the  oxidation  obtained  by  titrating  with  perman- 
ganate the  ferrocyanide  formed. 

According  to  this  method  an  excess  of  at  least  four  to  six  times 
the  theoretical  amount  of  potassium  ferricyanide  and  40  cm.3  to  50 
cm.3  of  a  2-normal  solution  of  sodium  hydroxide  are  added  to  the 
solution  containing  the  chromium  to  insure  complete  oxidation. 
The  oxidized  chromium  is  removed  by  precipitation  as  barium 
chromate  by  means  of  barium  hydroxide  in  solution,  and  fil- 
tration. The  filtrate  is  acidified  with-  hydrochloric  acid  and 
titrated  with  permanganate  according  to  Bollenbach 's  modifica- 
tion of  De  Haen's  method,  f  designed  to  overcome  the  difficulty 
involved  in  obtaining  an  exact  end  point  in  the  titration  of  large 
amounts  of  ferrocyanide,  owing,  as  was  first  pointed  out  by 
Griitzner,  \  to  the  formation  of  a  precipitate  of  K2MnFeC6N6 
during  the  titration.  This  modification  consists  in  adding  an 
excess  of  permanganate  to  the  solution  and,  after  the  precipitate 
has  cleared  up,  titrating  back  the  excess  of  permanganate  with 
n/2O  potassium  ferrocyanide  in  the  presence  of  a  trace  of  a  ferric 
salt,  the  formation  of  a  permanent  green  coloration  due  to  the 
ferric  ferrocyanide  indicating  the  end  point. 

According  to  the  experience  of  Palmer  §  the  proportions  of 
ferricyanide  and  alkali  hydroxide  prescribed  above  are  insuf- 
ficient to  bring  about  complete  oxidation  of  the  chromium  in 
chromic  condition  to  that  of  chromic  acid,  according  to  the 
equation 

Cr203  +  6  K3FeC6N6  +  6  KOH  =  2  CrO3  +  6  K4FeC6N6  +  3  H2(X 

*  Zeit.  anorg.  Chem.,  Ix,  446. 

t  Zeit.  anal.  Chem.,  xlvii,  687. 

J  Chem.  Centralblatt,  1902,  i,  500. 

§  Howard  E.  Palmer,  Am.  Jour.  Sci.,  [4],  xxx,  141. 


414 


METHODS  IN  CHEMICAL  ANALYSIS 


But  by  using  about  fifteen  times  the  theoretical  proportion  of 
potassium  ferricyanide  and  a  rather  strong  solution  of  potassium 
hydroxide  in  a  total  volume  of  solution  amounting  to  100  cm.3 
or  125  cm.3  results  are  obtained  in  accord  with  the  theory. 

Palmer  tested  the  process  upon  a  chromic  salt  made  by  treat- 
ing portions  of  a  standardized  solution  of  potassium  chromate, 
made  slightly  acid  with  hydrochloric  acid,  with  a  current  of 
^sulphur  dioxide  until  the  clear  green  color  indicated  complete 
reduction  of  the  chromic  acid  to  the  condition  of  chromic  oxide, 
Cr2O3.  The  sulphur  dioxide  was  then  expelled  by  boiling  the 
solution  in  a  current  of  carbon  dioxide.  To  determine  the  chro- 
mium in  such  a  solution  the  following  procedure  proved  effective. 

To  the  cold  solution  of  the  chromic  salt  are  added  15  to  20 
times  the  theoretical  amount  of  potassium  ferricyanide  and  potas- 
sium hydroxide  in  rather  strong  solution,  with  care  to  make  the 
final  volume  100  cm.3  to  125  cm.3.  Barium  hydroxide  is  added 
to  precipitate  the  chromate  formed  and  the  insoluble  barium 
chromate  is  removed  by  filtration.  The  solution  is  acidified  with 
hydrochloric  acid  and  treated  with  measured  permanganate  in 
excess,  and  the  excess  of  permanganate  is  determined  by  titration 
to  permanent  green  with  potassium  ferrocyanide  in  presence  of  a 
ferric  salt,  with  due  correction  for  the  amount  of  permanganate 
taken  up  by  the  same  amount  of  the  ferricyanide  alone. 

Experimental  results  are  given  in  the  table. 

Estimation  of  Chromium  in  the  Chromic  Condition. 


CrO3  taken. 

K3FeC6N6  used. 

KOH  used. 

CrO3  found. 

Error. 

grm. 

grm. 

grm. 

grm. 

grm. 

O.IOIO 
O.IOIO 
O.IOIO 
O.IOIO 

6 
8 
8 
8 

12 

16 
16 

12 

0.0979 
0.0981 
o  .  0989 

0.0997 

-0.0031 
—  0.0029 
—  O.OO2I 
—  0.0013 

MOLYBDENUM. 

The  Gravimetric  Estimation  of  Molybdic  Acid  by  Liberation  of 
Iodine  and  Absorption  of  that  Element  by  Silver. 

^'  When  a  soluble  molybdate  is  added  to  an  excess  of  potassium 
iodide  made  acid  with  hydrochloric  acid  and  shaken  with  elec- 
trolytically  prepared  silver  under  an*  atmosphere  of  hydrogen,* 
iodine  equivalent  to  the  molybdic  acid  is  set  free  and  then  is 

*  See  page  27. 


MOLYBDENUM 


415 


absorbed  by  the  silver.  From  the  increase  in  weight  of  the  silver 
the  amount  of  molybdenum  trioxide  was  calculated  on  the 
assumption  that  one  molecule  of  molybdenum  trioxide  liberates 
one  atom  of  iodine  according  to  the  following  equation: 

2  MoO3  +  4  KI  +  4  HC1  =  2  MoO2I  +  4  KC1  +  2  H2O  +  I2. 

Results  obtained  by  Perkins*  in  the  treatment  of  ammo- 
nium molybdate,  the  composition  of  which  had  been  determined 
by  fusion  with  sodium  tungstate  containing  a  slight  excess  of 
tungstic  acid,  are  given  in  the  table. 

Absorption  of  Iodine  by  Silver. 


Ag  taken, 
grm. 

MoO3  taken, 
grm. 

I2  found, 
grm. 

Calculated  MoO3. 
grm. 

Error, 
grm. 

2.0002 

o.  2127 

0.1869 

O.  2I2O 

—  O.OOO7 

2.OOO6 

o.  2127 

o.  1874 

o.  2126 

—  O.OOOI 

2.OOI2 

o.  2127 

o.  1870 

O.2I2I 

—  O.OOO6 

2.0048 

o.  2127 

o.  1876 

o.  2128 

+O.OOOI 

2.OOOO 

0.2540 

0.2242 

0.2543 

+0.0003 

2.0004 

o.  2909 

0.2571 

0.2916 

+0.0007 

The  lodometric  Estimation  of  Molybdic  Acid. 

The  Digestion  Mauro  and  Danesi  f  have  made  use  of  the  reac- 
Method.  tjon  wnich  takes  place  between  hydrochloric  acid, 

potassium  iodide  and  a  soluble  molybdate  to  determine  the 
amount  of  molybdic  acid  from  the  amount  of  iodine  set  free  in 
accordance  with  the  reaction 

2  MoO3  +  4  HI  =  2  MoO2I  +  I2  +  2  H2O. 

The  best  results  are  obtained  by  acting  upon  a  soluble  molybdate 
containing  from  o.i  to  0.5  grm.  of  molybdic  acid,  with  1.5  grm. 
of  potassium  iodide  in  1.5  cm3,  of  water  and  2.5cm3.  of  strong 
hydrochloric  acid  in  an  atmosphere  of  carbon  dioxide,  the  whole 
being  heated  an  hour  and  a  half  in  a  sealed  tube.  The  authors 
point  out  that  with  prolonged  heating  the  action  proceeds  a 
little  further,  and  in  the  cold,  under  conditions  otherwise 
similar,  not  quite  so  far  as  the  theory  of  the  equation  would 
indicate. 

*  Claude  C.  Perkins,  Am.  Jour.  Sci.,  [4],  xxix,  540. 
t  Zeit.  anal.,  Chem.,  xx,  567. 


41 6  METHODS  IN  CHEMICAL  ANALYSIS 

In  studying  this  method  Gooch  and  Fairbanks  *  have  obtained 
variable  results  fairly  in  accord  with  the  theory  of  the  reduction 
when  the  digestion  is  made  in  sealed  tubes  at  extremely  small 
volume,  and  with  small  amounts  of  the  molybdate,  under  the 
exact  conditions  indicated,  but  widely  deficient  for  larger  amounts, 
for  larger  volumes,  and  for  digestions  in  the  cold.  This  behavior 
indicates,  of  course,  a  tendency  on  the  part  of  the  iodine  to  re- 
verse the  action,  and  the  obvious  remedy  for  the  reversal  should 
be  found  in  the  removal  of  the  iodine  as  it  is  set  free.  Fried- 
heim  and  Euler  f  accomplish  this  by  a  process  of  distillation  in 
the  Bunsen  apparatus,  collecting  and  determining  the  iodine  in 
the  distillate. 

Distillation  The  experience  of  Gooch  and  Fairbanks  t  confirms 

process.  'm  generai  the  utility  of  the  distillation  process  pro- 

vided that  conditions  are  exactly  defined.  It  is  not  sufficient 
to  say  that  the  boiling  should  be  stopped  when  a  clear  green  color 
appears  and  when  the  steam  is  no  longer  colored  by  iodine ;  for 
the  green  color  comes  very  gradually,  and  iodine  remains  in  the 
residue  after  the  green  color  has  developed  distinctly.  It  is  safer 
and  more  convenient  to  start  the  distillation  with  a  definite  volume 
of  liquid  and  boil  until  the  volume  is  reduced  to  a  definite  point. 
If  the  initial  volume  is  made  about  40  cm3.,  no  iodine  remains  in 
the  flask  after  the  liquid  has  been  boiled  down  to  25  cm3.,  and 
at  that  degree  of  concentration  the  molybdic  acid  shows  the  theo- 
retical reduction;  but  if  the  concentration  is  pushed  beyond  this 
point,  a  tendency  to  further  reduction  of  the  molybdic  acid 
becomes  evident. 

It  is  necessary  to  carry  on  the  distillation  in  an  atmosphere 
of  carbon  dioxide,  inasmuch  as  the  hydriodic  acid  freed  by  the 
action  of  hydrochloric  acid  upon  the  potassium  iodide  is  decom- 
posed by  distillation  in  contact  with  air,  with  liberation  of  iodine. 
As  even  a  trace  of  oxygen  will  immediately  set  free  iodine  from 
boiling  hydriodic  acid,  the  carbon  dioxide  is  best  evolved  from 
boiled  marble  by  the  action  of  boiled  acid  to  which  a  little  cu- 
prous chloride  has  been  added,  and  finally  passed  through  solu- 
tions of  iodine  and  potassium  iodide  to  free  it  from  any  reducing 
substance. 

*  F.  A.  Gooch  and  Charlotte  Fairbanks,  Am.  Jour.  Sci.,  [4],  ii,  156. 
t  Ber.  Dtsch.  Chem.  Ges.,  xxviii,  2066. 
t  Loc.  cit. 


MOLYBDENUM 


417 


A  convenient  apparatus,  constructed  with  sealed  and  ground 
joints  exclusively,  is  shown  in  the  accompanying  figure.  The 
distillation  takes  place  in  the  first  flask, 
and  the  iodine  collects  in  the  second  flask 
and  trap,  which  hold  a  solution  of  potas- 
sium iodide  kept  cool  by  immersion  of 
the  flask  in  cold  water. 

In  carrying  out  a  determination  with 
this  apparatus,  the  purified  carbon  di- 
oxide is  first  passed  for  some  minutes 
and  the  stopcock  in  the  funnel  is  then 
closed.  The  molybdate  dissolved  in  10 
cm.3  of  boiled  water  is  put  in  the  funnel 
and  almost  all  of  it  is  allowed  to  run 
into  the  first  flask.  It  is  necessary  that 
a  few  drops  be  left  in  the  funnel  so  that 
the  liquid  to  follow  may  not  carry  down  bubbles  of  air.  Potas- 
sium iodide,  never  exceeding  the  theoretical  requirement  by 
more  than  0.5  grm.,  is  introduced  similarly  in  10  cm.3  of  boiled 
water,  followed  by  20  cm.3  of  concentrated  hydrochloric  acid 
(sp.  gr.  i. 20).  Before  the  acid  is  allowed  to  run  in  entirely,  the 
funnel  is  again  filled  with  carbon  dioxide  and  finally  left  con- 
nected with  the  generator  so  that  carbon  dioxide  may  be  passed 
into  the  apparatus  at  pleasure.  The  liquid  in  the  first  flask  is 
boiled  until  the  volume  of  liquid  has  decreased  to  25  cm.3, 
indicated  by  a  mark.  The  iodine  collected  in  the  second  flask 

Distillation  Process. 


Ffg.  27. 


MoO3  taken  as 
ammonium  molybdate. 

grm. 

KI. 

grm. 

MoO3  found, 
grm. 

Error, 
grm. 

0.2585 

0-5 

0.2580 

—  0.0005 

0.2995 

0.5 

O.  2991 

—  0.0004 

0.2524 

0-5 

0.2513 

—  O.OOII 

o  .  2446 

0-5 

0.2457 

+0.001  1 

0.2903 

0-5 

0.2914 

-j-o.oon 

0.2798 

0-5 

0.2808 

-j-o.ooio 

o.  2656 

MoO3  dissolved  in 

0-5 

0.2663 

+0.0007 

NaOH. 

grm. 

0.2273 

o-S 

0.2281 

+o  .  0008 

0.2052 

0.5 

o.  2062 

+0.0010 

0-3474 

0.5 

0.3467 

—0.0007 

4i8 


METHODS  IN  CHEMICAL  ANALYSIS 


and  trap  is  titrated  with  standardized  thiosulphate.  Results  of 
test  experiments  made  according  to  this  procedure  are  given  in 
the  table. 

Further  study  of  this  process  has  been  made  by  Gooch  and 
Norton,*  with  a  view  to  testing  further  the  reliability  of  the 
process  and  to  determining  the  progress  of  the  reduction  of 
molybdic  acid  as  the  concentration  proceeds. 

The  distillation  apparatus  employed  in  this  work  was  con- 
structed with  sealed  or  ground  joints  of  glass  wherever  contact 
with  iodine  was  a  possibility.  It  was  made  by  sealing  together 
a  separating  funnel  A,  a  100  cm.3  Voit  flask  B,  a  Drexel  wash 
bottle  C,  and  a  bulbed  trap  g,  as  shown  in  Figure  3-f  Upon  the 
side  of  the  distillation  flask  B  was  pasted  a  graduated  scale,  by 
means  of  which  the  volume  of  the  liquid  within  the  flask  might 
be  known  at  any  time.  Carbon  dioxide,  generated  in  a  Kipp 
apparatus  by  the  action  of  dilute  hydrochloric  acid  (carrying  in 
solution  cuprous  chloride  to  take  up  free  oxygen)  upon  marble 
previously  boiled  in  water,  was  passed  through  the  apparatus 
before  and  during  the  operation,  so  that  it  was  possible  to  inter- 
rupt the  process  of  boiling  at  any  point  of  concentration,  to 
remove  the  receiver  by  easy  manipulation,  to  replace  the  receiver, 
and  to  continue  the  distillation  without  danger  of  admitting  air 
to  the  distillation  flask. 

The  mean  error  of  the  indications  found  for  different  periods 
of  distillation  are  given  in  the  accompanying  statement. 


Concentration. 

Error. 

cm.J 

grrn. 

40  to  32 
40  to  25 
40  to  10 

-0.0045 
+0.0008 
+0.0014 

The  best  results,  showing  a  mean  error  of  -f  0.0008  grm. 
between  extremes  of  —  o.oooi  grm.  and  +0.0020  grm.,  were 
obtained  by  this  process  when  the  distillation  was  continued  until 
the  original  volume  of  40  cm.3  had  been  diminished  to  25  cm.3. 
Concentration  beyond  the  limit  of  25  cm.3  plainly  develops  a 

*  F.  A.  Gooch  and  John  T.  Norton,  Jr.,  Am.  Jour.  Sci.,  [4],  vi,  168. 
t  See  page  4. 


MOLYBDENUM  419 

tendency  toward  over-reduction,  especially  when  the  amount 
of  potassium  iodide  is  increased  beyond  about  0.5  grm.  in  excess 
of  that  theoretically  required. 

It  is  shown,  further,  that  the  precaution  of  conducting  the  oper- 
ation in  an  atmosphere  of  carbon  dioxide  does  not  eliminate  all 
chance  of  error  of  this  sort  unless  the  liquid  of  the  mixture  — 
the  hydrochloric  acid  as  well  as  the  water  —  is  free  from  air. 
Thus,  40  cm.3  of  unboiled  acid,  sp.  gr.  1.12,  introduced  enough 
air  into  the  apparatus  to  cause  an  error  of  0.0013  grm.  reckoned 
in  terms  of  molybdenum  trioxide,  while  the  iodine  set  free  by 
the  action  of  the  residual  acid  of  this  experiment  upon  another 
gram  of  potassium  iodide  introduced  without  admission  of  air 
corresponded  to  only  0.0002  grm.  in  terms  of  molybdenum  triox- 
ide. The  use  of  acid  of  sp.  gr.  i.i,  freshly  boiled  in  the  air, 
obviously  reduces  the  error  due  to  the  unboiled  acid,  but  even 
in  this  case  the  effect  of  dissolved  oxygen  is  not  wholly 
obviated. 

While  it  is  possible  to  determine  molybdic  acid  by  esti- 
mating the  iodine  set  free  upon  boiling  a  solution  of  that  acid 
in  hydrochloric  acid  to  which  potassium  iodide  has  been  added, 
the  operation  is  beset  with  difficulties.  If  due  attention  be  paid 
to  the  proportion  of  acid,  the  excess  of  potassium  iodide  over  the 
amount  theoretically  necessary  to  produce  the  ideal  reduction, 
and  the  final  degree  of  concentration  of  the  liquid,  the  change  of 
the  molybdic  acid  to  the  condition  of  oxidation  of  the  pentoxide, 
Mo2O5,  may,  as  has  been  shown,  be  closely  realized.  Even  when 
the  conditions  essential  to  the  ideal  reduction  obtain,  however, 
the  possibility  of  the  interaction  of  atmospheric  oxygen  with  the 
hydriodic  acid  produced  from  the  potassium  iodide  and  hydro- 
chloric acid  must  be  guarded  against  by  conducting  the  oper- 
ation in  an  atmosphere  of  carbon  dioxide.  Furthermore,  the 
hydrochloric  acid  employed  must  be  freed,  so  far  as  may  be,  from 
dissolved  oxygen  by  previous  boiling.  The  operation  is  capable 
of  yielding  accurate  indications  (unless  by  a  chance  combination 
of  opposite  errors)  only  when  the  precautions  mentioned  are 
scrupulously  observed. 

On  the  other  hand,  the  essential  principles  of  the  reaction 
admit  of  very  simple  application  when  the  residue  of  the  distilla- 
tion, instead  of  the  iodine  distilled,  is  made  the  object  of  the 
analytical  determination. 


420  METHODS  IN  CHEMICAL  ANALYSIS 

Reoridationof  According  to  the  method  developed  by  Gooch  and 
the  Residue  by  Fairbanks  *  the  mixture  of  molybdic  acid,  potassium 
iodide,  and  hydrochloric  acid  is  boiled  between  de- 
fined limits  of  volume  in  an  Erlenmeyer  flask  simply  trapped  to 
prevent  mechanical  loss  during  the  boiling.  The  residue  after 
the  removal  of  all  free  iodine  is  treated  with  tartaric  acid  to 
prevent  subsequent  precipitation,  neutralized  with  an  alkali 
bicarbonate,  and  reoxidized  by  a  measured  excess  of  standard 
iodine,  this  excess  of  iodine  being  finally  determined  by  titra- 
tion  with  standard  arsenite.  The  difference  between  the  iodine 
value  of  the  arsenite  used  and  the  amount  of  standard  iodine 
employed  measures  the  molybdic  acid.  The  action  of  atmos- 
pheric or  dissolved  oxygen  obviously  plays  no  important  part 
in  the  operation,  provided  all  uncombined  iodine,  however  pro- 
duced, is  finally  boiled  out.  At  the  outset  iodine  is  liberated  by 
the  air  present  as  well  as  by  the  molybdic  acid,  but  as  the 
boiling  continues  the  hydriodic  acid  diminishes  in  strength, 
the  iodine  is  driven  from  the  flask  kept  filled  with  steam,  and 
if  the  conditions  of  the  operation  have  been  properly  adjusted 
the  molybdic  acid  undergoes  the  ideal  reduction.  When  this 
point  is  reached  dilution  with  cold  water  obviates  danger  of 
immediate  oxidation  and  gives  opportunity  for  continuing  the 
treatment  referred  to  above. 

According  to  this  method,  the  soluble  molybdate  in  amount  not 
exceeding  the  equivalent  of  0.5  grm.  of  MoOs,  and  at  least  20  cm.3 
of  hydrochloric  acid  (sp.  gr.  1.20)  with  from  0.2  grm.  to  0.6  grm. 
of  potassium  iodide,  according  to  the  amount  of  molybdate  used, 
are  put  into  a  150  cm.3  flask  which  is  then  trapped  loosely  by 
a  short  bulbed  tube  hung  in  the  neck,  as  shown  in  Fig.  6.f 
The  solution  is  boiled  until  the  original  volume  of  40  cm.3  to 
60  cm.3  has  been  reduced  to  exactly  25  cm.3,  as  determined  by  a 
mark  upon  the  flask.  The  residue  is  diluted  immediately  to 
a  volume  of  125  cm.3,  cooled,  and  transferred  to  a  reaction 
bottle  of  the  form  shown  in  Fig.  y,J  the  Will  and  Varrentrapp 
absorption  tube  being  charged  with  a  solution  of  potassium 
iodide  to  catch  any  traces  of  iodine  thrown  off  mechanically 
during  the  process.  Through  the  stoppered  funnel  is  added  in 

*  F.  A.  Gooch  and  Charlotte  Fairbanks,  Am.  Jour.  Sci.,  [4],  ii,  160. 
t  See  page  6. 
t  See  page  6. 


MOLYBDENUM 


421 


solution  i  grm.  of  tartaric  acid  to  prevent  precipitation  during 
the  subsequent  neutralization.  The  free  acid  is  neutralized  by 
introduction  of  acid  alkali  carbonate  (or  of  alkali  hydroxide  to 
partial  neutralization,  followed  by  the  acid  carbonate)  and  a 
measured  amount  of  n/io  iodine  in  excess  run  in.  The  iodine 
color  begins  to  fade  perceptibly  within  fifteen  minutes,  but  for 
complete  oxidation  the  mixture,  protected  from  direct  sunlight, 
should  be  set  aside  for  an  hour  and  a  half  or  two  hours.  The 
iodine  remaining  is  then  titrated  by  standard  arsenite. 

At  the  end  of  this  operation  it  is  wise  to  acidulate  the  solution 
with  dilute  hydrochloric  acid  and  determine  by  titration  with 
sodium  thiosulphate  any  slight  amount  of  iodine  which  may 
have  taken  the  form  of  iodate  in  the  long  digestion. 

Results  of  experimental  tests  of  this  procedure  are  given  in  the 

table. 

Reoxidation  of  Residue  by  Iodine. 


MoO3  taken  as 
ammonium  molybdate. 

grm. 

KI. 

grm. 

MoO-i  found, 
grm. 

Error, 
grm. 

Neutralized  by  NaHCO3 


0.1517 

o-S 

0.1517 

o.oooo 

0.2530 

0-5 

0-2537 

+0.0007 

o.  1636 

o-S 

0.1637 

+0.0001 

0.1702 

0.5 

0.1702 

0.0000 

0.1520 

0.5 

0.1518 

—  O.OOO2 

0.1642 

o-S 

0.1652 

+O.OOIO 

0.4560 

0-75 

o  .  4560 

o.oooo 

0.1690 

o-5 

0.1683 

—0.0007 

Partially  neutralized  by  NaOH:  fully  neutralized  by  NaHCO3. 


0.0507 

o-5 

0.0519 

+6.OOI2 

0.1663 

o-5 

0.1666 

+0.0003 

O.OIOI 

0.5 

0.0095 

—0.0006 

0.1639 

o.S 

0.1632 

—0.0007 

0.1636 

o-S 

0.1625 

—  o.oon 

0.0507 

o-5 

0.0510 

+o  .  0003 

0.1685 

0-5 

0.1683 

—  O.OOO2 

0.1514 

o-5 

0.1512 

—  O.OOO2 

o.  1649 

o-5 

0.1646 

—  0.0003 

of  method  just  described,  according  to  which  a 

the  Residue  by    soluble  molybdate  is  boiled  with  hydrochloric  acid 
langanate.    an(j  a  smajj  excess  of  potassium  iodide  to  a  definite 
concentration,  the  residue  neutralized  with  an  acid  alkali  carbon- 
ate, and  the  reduced  molybdic  salt  reoxidized  by  standard  iodine, 


422  METHODS  IN  CHEMICAL  ANALYSIS 

affords  an  accurate  determination  of  the  molybdate  but  is  some- 
what tedious  on  account  of  the  delay  necessary  before  the  final 
titration  by  standard  arsenite  may  be  effected. 

Gooch  and  Pulman  *  have  succeeded  in  obviating  the  difficulty 
of  long  delay  by  substituting  potassium  permanganate  for  iodine 
in  the  process  of  reoxidation. 

When  an  excess  of  potassium  permanganate  is  added  to  a 
solution  containing  hydrochloric  and  hydriodic  acids  with  molyb- 
denum in  a  condition  of  oxidation  corresponding  to  the  pentox- 
ide,  several  different  effects  of  oxidation  may  be  expected. 
There  is  the  immediate  liberation  of  iodine  from  the  hydriodic 
acid,  the  production  of  iodic  acid,  some  slight  tendency  to  lib- 
erate chlorine  from  the  hydrochloric  acid,  the  formation  of 
representatives  of  the  higher  oxides  of  manganese,  and,  lastly, 
the  reproduction  of  the  molybdic  acid,  which  is  the  object  of  the 
operation.  All  other  effects  than  the  last  must  be  prevented  or 
recorded  in  order  that  the  estimation  of  the  molybdic  acid  may 
be  accomplished.  Of  the  secondary  actions,  the  liberation  of 
chlorine  may  be  prevented  by  adding  a  manganous  salt,  accord- 
ing to  the  well-known  proposals  of  Kessler  f  and  Zimmerman;  t 
the  iodine  set  free  may  be  converted  to  hydriodic  acid  again  by 
the  introduction  into  the  acid  solution  of  a  sufficient  excess  of  a 
standard  arsenite  solution;  the  iodic  acid  may  be  reduced  in 
the  acid  solution  in  the  same  manner,  as  was  found  by  experi- 
ment, since  the  value  obtained  for  a  solution  of  iodic  acid  by 
adding  to  it  in  presence  of  sulphuric  acid  a  measured  amount 
of  standard  arsenite,  making  alkaline,  and  titrating  back  with 
iodine,  proved  to  be  the  same  as  that  found  by  acting  directly 
upon  the  iodic  acid  with  an  excess  of  potassium  iodide  in  pres- 
ence of  dilute  sulphuric  acid  and  determining  by  standard  arse- 
nite the  iodine  liberated.  The  higher  oxides  of  manganese  are 
likewise  reduced  in  the  acid  solution  by  the  arsenious  acid  of  the 
standard  arsenite. 

Whatever  excess  of  arsenious  acid  is  left  over  after  the  reac- 
tions described  may,  obviously,  be  determined  by  neutralizing 
the  solution  with  an  alkali  bicarbonate  and  titrating  with 
iodine. 

*  F.  A.  Gooch  and  O.  S.  Pulman,  Jr.,  Am.  Jour.  Sci.,  [4],  xii,  449. 
t  Ann.  Phys.,  cxviii.  48;   cxix,  225-226. 
t  Ann.  Chem.,  ccxiii,  302. 


MOLYBDENUM  423 

Every  difficulty  introduced  by  the  secondary  oxidations  may 
be  overcome  by  adopting  the  following  procedure:  the  addition 
of  manganous  sulphate  to  the  reduced,  diluted  and  cooled  solu- 
tion ;  the  introduction  of  measured  standard  potassium  perman- 
ganate until  its  characteristic  color  is  evident;  the  treatment  of 
the  solution  with  a  measured  amount  of  standard  arsenite  in 
known  excess,  to  destroy  the  excess  of  permanganate ;  the  intro- 
duction of  tartaric  acid  to  prevent  subsequent  precipitation; 
neutralization  by  acid  alkali  carbonate ;  and  titration  of  the  excess 
of  arsenite  by  iodine. 

The  operation  requires  the  use  of  a  standard  solution  of  arse- 
nite, easily  made  with  accuracy,  a  solution  of  iodine  in  potassium 
iodide  standardized  directly  against  the  arsenite  in  the  usual 
manner,  and  a  solution  of  potassium  permanganate  the  value 
of  which  in  terms  of  the  arsenite  is  found  by  bleaching  a  meas- 
ured portion  with  an  excess  of  arsenite  and  titrating  back  with 
iodine  the  arsenite  remaining.  The  value  (in  terms  of  molybdic 
acid)  of  the  permanganate  used,  diminished  by  that  of  the  arse- 
nite and  increased  by  that  of  the  iodine,  gives  the  amount  of 
molybdic  acid  present. 

According  to  this  procedure  the  soluble  molybdate  in  amount 
not  exceeding  the  equivalent  of  0.5  grm.  of  MoO3,  at  least 
20  cm.3  of  hydrochloric  acid  (sp.  gr.  1.20),  with  0.2  grm.  to  O.6 
grm.  of  potassium  iodide,  according  to  the  amount  of  molybdate 
used,  are  put  into  a  150  cm.3  flask  trapped  by  a  short  bulbed  tube 
hung  loosely  in  the  neck,  as  shown  in  Fig.  6  on  p.  6.  The  solu- 
tion, having  a  volume  of  40  cm.3  to  60  cm.3  originally,  is  boiled 
to  a  volume  of  25  cm.3,  indicated  by  a  mark  upon  the  flask. 
The  residue  is  diluted  immediately  to  a  volume  of  125  cm.3, 
cooled,  and  transferred  to  a  bottle  fitted  with  a  stoppered  funnel 
and  trap,  as  shown  in  Fig.  6,  p.  6.  Through  the  stoppered 
funnel  are  added  0.5  grm.  of  manganese  sulphate  and  a  meas- 
ured amount  of  n/io  potassium  permanganate  to  the  character- 
istic coloration.  A  measured  amount  of  standard  arsenite  is 
then  added,  enough  to  correspond  approximately  to  the  perman- 
ganate used,  experience  having  shown  that  this  amount  of  arse- 
nite is  sufficient  to  reduce  with  readiness  in  the  acid  solution  the 
iodic  acid,  the  permanganate,  the  higher  oxides  of  manganese,  and 
nearly  all  the  free  iodine,  the  remainder  of  the  last  being  taken 
up  immediately  after  the  subsequent  neutralization.  Next,  a 


424 


METHODS  IN  CHEMICAL  ANALYSIS 


solution  of  about  3  grm.  of  tartaric  acid  is  run  in,  and  the  free 
acid  of  the  solution,  is  neutralized  by  acid  potassium  carbonate. 
Finally,  the  liquid  adhering  to  stopper  and  tubes  is  washed  off 
into  the  bottle,  the  contents  of  the  trap  are  added,  and  the  residual 
arsenite  is  titrated  by  standard  iodine,  using  the  starch  indicator. 

When  many  determinations  are  to  be  made  the  process  yields 
results  with  rapidity.  Many  operations  may  be  started  suc- 
cessively in  the  Erlenmeyer  flasks,  and  one  neutralization  bottle 
may  serve  for  the  treatment  of  all  the  reduced  residues  as  they 
come  along. 

The  results  of  experiments  are  given  in  the  following  table. 
Reoxidation  of  Residue  by  Permanganate. 


Weight  of  MoO3  taken 
as  ammonium 
molybdate. 

grin. 

Weight  of  KI  used, 
grm. 

Weight  of  Mo03  found, 
grm. 

Error, 
grm. 

0.0423 
0.0429 
0.0420 
0.0827 
0.0837 
0.0826 
0.2465 
0.2481 
0.2470 

0.2 
0.2 
O.  2 
0-3 

o-3S 
o-S 
0.6 
0.6 
0.6 

o  .  0430 
0.0435 
0.0427 
0.0829 
0.0838 
0.0832 
o  .  2460 
o  .  2469 

0.2465 

+0  .  0007 
+0.0006 
+0.0007 
+O.OOO2 
+O.OOOI 

+o  .  0006 

—  0.0005 
—  O.OOI2 
—  O.OOO5 

The  Estimation  of  Molybdic  Acid  reduced  in  the  Jones  Redactor. 

Opinions  have  differed  in  regard  to  the  degree  of  reduction 
obtained  when  molybdic  trioxide,  in  sulphuric  acid  solution,  is 
passed  through  the  column  of  amalgamated  zinc  as  applied  in 
the  Jones  reductor.  Jones,*  who  first  determined  molybdenum 
by  this  method,  considered  that  the  reduction  goes  to  the  condi- 
tion represented  by  the  formula  Moi2Oi9,  the  same  degree  of 
reduction  that  Wernke  f  obtained  with  zinc  and  sulphuric  acid 
in  a  closed  flask.  Doolittle  %  and  Eavenson  found  that,  by 
varying  the  strength  of  the  acid  and  the  speed  at  which  the 
molybdenum  was  passed  through,  different  degrees  of  reduction 
might  be  obtained,  but  none  were  lower  than  that  represented 
hy  the  formula  Moi2Oi9.  Blair  §  and  Whitfield  were  unable  to 

*  Am.  Inst.  Min.  Eng.,  xviii,  705. 

f  Zeit.  anal.  Chem.,  xiv.,  i. 

J  Jour.  Am.  Chem.  Soc.,  xvi,  234. 

§  Ibid.,  xvii,  747. 


MOLYBDENUM 


425 


press  the  reduction  below  the  condition  represented  by  the  sym- 
bol Mo24O37.  Miller*  and  Frank  in  repeating  the  experiments 
of  Blair  and  Whitfield  obtained  in  general  the  same  results, 
though  by  taking  extraordinary  precautions  they  were  able  to 
get  a  reduction  to  a  little  below  the  midway  point  between  the 
conditions  represented  by  the  symbols  Mc^Osv  and  Mo2O3. 
W.  A.  Noyes  and  Frohman,f  by  taking  pains  to  replace  the  air 
in  the  reductor  flask  by  carbon  dioxide,  were  able  to  obtain  a 
reduction  to  the  form  of  Mo2O3. 

Noting  the  ease  with  which  the  reduced  molybdenum  solution 
is  oxidized  by  the  air,  Randall  t  has  investigated  the  possibility 
of  charging  the  reductor  flask  with  an  oxidizer  unaffected  by 
air,  to  anticipate  the  oxidizing  effects  of  the  air  as  the  reduced 
molybdenum  compound  comes  through  the  reductor,  and  to  reg- 
ister the  oxidation.  In  preliminary  experiments  an  excess  of  a 
standard  solution  of  potassium  permanganate  was  used  in  the 
receiver.  But  it  was  found,  in  blank  tests,  that  somewhat  more 
than  the  theoretical  amount  of  permanganate  was  used  up,  due 
either  to  impurities  in  the  zinc,  to  small  particles  of  zinc  which 
had  worked  through  the  asbestos  at  the  bottom  of  the  reductor, 
or  possibly  to  the  hydrogen  formed  in  the  reductor.  That  the 
last  mentioned  possibility  is  a  sufficient  cause  of  the  effects 
obtained  was  shown  by  passing  hydrogen,  formed  by  the  action 
of  hydrochloric  acid  on  zinc  in  a  Kipp  generator  and  washed  with 
water  and  caustic  potash,  through  23  cm.3  of  standard  perman- 
ganate diluted  with  300  cm.3  of  hot  dilute  sulphuric  acid  (2.5  per 
cent) ,  adding  20  cm.3  of  a  solution  of  standard  ferrous  sulphate 
and  titrating  back  with  permanganate.  The  results  obtained 
by  such  action  of  hydrogen  during  fifteen  minutes  are  shown  in 
the  accompanying  statement: 

Effect  of  Hydrogen. 


FeS04. 

KMnO4  required. 

KMnO4  theory. 

Reduced  by  hydrogen. 

cm.8 

cm.» 

cm.3 

cm.8 

20 
2O 
20 

24.6 

24-7 
23-6 

23.1 
23.1 
23.1 

i.  5 

1.6 

o-S 

*  Jour.  Am.  Chem.  Soc.,  xxv,  919. 

t  Ibid.,  xvi,  553. 

$  D.  L.  Randall,  Am.  Jour.  Sci.,  [4],  xxiv,  313. 


426 


METHODS  IN  CHEMICAL  ANALYSIS 


The  use  of  ferric  alum  *  in  the  receiving  flask,  with  phosphoric 
acid  to  decolorize  it,f  Randall  found  to  be  unobjectionable  and 
effective.  According  to  the  procedure  laid  down  by  Randall, 
the  receiving  flask  J  is  charged  with  20  cm.3  to  30  cm.3  of  a  solution 
prepared  by  dissolving  100  grm.  of  ferric  alum  in  a  liter  of  water  j 
phosphoric  acid  (4  cm.3  of  the  sirupy  acid)  is  added  to  the  solu- 
tion in  the  receiver;  and  through  the  36  cm.  column  of  amalga- 
mated zinc  in  the  reductor  are  passed  in  succession  100  cm.3  of 
hot  dilute  sulphuric  acid  (2.5  per  cent),  the  molybdic  acid  in 
the  form  of  ammonium  molybdate  dissolved  in  10  cm.3  of  water 
and  acidified  with  100  cm.3  of  the  hot  dilute  acid  (2.5  per 
cent),  200  cm.3  of  the  hot  dilute  acid,  and  finally  100  cm.3  of 
hot  water.  The  molybdenum  salt  is  green  as  it  passes  through 
the  lower  part  of  the  reductor,  but  on  coming  in  contact 
with  the  ferric  salt  it  is  changed  to  a  bright  red.  The  solution 
is  titrated  while  still  hot  with  approximately  tenth  normal 
permanganate. 

The  results  in  the  following  table  are  calculated  on  the  assump- 
tion that  the  molybdic  acid  is  reduced  to  the  form  of  Mo2O3  and 
the  close  agreement  with  theory  indicates  that  the  reduction  does 
not  stop  at  an  intermediate  point. 

Collection  of  Mo2Oa  in  Ferric  Sulphate:    Titration  of  Ferrous  Salt  by  Perman- 
ganate. 


Ammonium 
molybdate 
taken. 

Iron 
solution. 

H3P04. 

KMnO4 
used. 

MoO3. 

Error. 

Found. 

Theory. 

grm. 

cm.' 

cm.3 

cm.* 

grm. 

grm. 

grill. 

O.2OOO 

2O 

4 

30.38 

0.1628 

0.1631 

—0.0003 

O  .  2OOO 

2O 

4 

30-45 

0.1632 

0.1631 

+  O.OOOI 

O.  2OOO 

2O 

4 

30.33 

0.1626 

0.1631 

—0.0005 

0.3000 

30 

4 

45-6S 

0.2447 

o  .  2447 

0.0000 

0.3000 

30 

4 

45.80 

0.2455 

0.2447 

+0.0008 

0.3000 

30 

4 

45-73 

0.2451 

0.2447 

+0.0004 

0.3000 

30 

4 

45-76 

0-2453 

0.2447 

+0.0006 

0.3000 

30 

4 

45-68 

0.2448 

0.2447 

+  0.0001 

0.3000 

30 

4 

45-64 

o  .  2446 

0.2447 

—  O.OOOI 

0.3000 

30 

4 

45-71 

0.2449 

0.2447 

+O  .  OO02 

*  Professor  Henry  Fay  kindly  gives  the  information  that  this  mode  of 
treating  reduced  molybdic  oxide  was  first  worked  out  many  years  ago  by  Dr. 
C.  B.  Dudley,  though  never  published. 

t  C.  Reinhardt,  Chem.  Ztg.,  13,  33. 

t  See  page  347. 


MOLYBDENUM 


427 


The  Determination  of  Molybdic  Acid  and  Vanadic  Acid  by  Reduc- 
tions and  Oxidations. 

Molybdic  acid  and  vanadic  acid  are  reduced  in  a  perfectly 
definite  manner  by  a  column  of  amalgamated  zinc,  and  each  by 
itself,  or  both  together,  may  be  estimated  by  titration  with 
potassium  permanganate  if  the  receiver  be  charged  with  a  solu- 
tion of  ferric  alum  to  anticipate  possible  oxidizing  action  of  the 
air.*  On  the  other  hand,  of  these  two  acids  only  vanadic  acid  is 
easily  reduced  in  solution  by  sulphur  dioxide.  These  are  char- 
acteristics which  suggested  to  Edgar  f  the  search  for  conditions 
under  which  molybdic  acid  would  escape  all  action  by  sulphur 
dioxide,  and  the  development  of  a  method  for  the  determination 
of  the  two  acids. 

Action  of  Sulphur  Dioxide  on  Molybdic  Acid. 


Volume  of 
solution. 

Mo03. 

H2SO4  (sp.  gr.  1.84). 

Time  of 
treatment 
with  SO2. 

KMn04 
w/ioXi.oo4. 

Color  of  solution. 

cm.3 

grm. 

cm.8 

min. 

cm.3 

25.0 
3S-° 

O.  2OO 
O.  2OO 

Faintly  acid. 
o-5 

IO 
IO 

0-15 
0.05 

Light  blue. 
Faint  blue. 

50.0 

O.  2OO 

i  .0 

IO 

o.o 

Colorless. 

75-o 

O.  20O 

2.O 

IO 

o.o 

Colorless. 

50.0 

O.  2OO 

2.O 

30 

0.0 

Colorless. 

25.0 

O.  200 

S-° 

IO 

0.0 

Colorless. 

25.0 

O.400 

S-o 

10 

0.0 

Colorless. 

50.0 

O.40O 

5-° 

IO 

o.o 

Colorless. 

50.0 

O.40O 

IO.O 

IO 

o.o 

Colorless. 

50.0 

0.400 

15-0 

10 

0.0 

Colorless. 

To  determine  the  conditions  under  which  molybdic  acid  is 
unaffected  by  sulphur  dioxide,  experiments  were  made  in  which 
solutions  of  molybdic  acid  of  varying  concentrations,  acidified 
with  varying  amounts  of  sulphuric  acid,  were  heated  to  boiling 
and  treated  with  a  current  of  sulphur  dioxide  for  varying  lengths 
of  time.  The  excess  of  sulphur  dioxide  was  then  removed  by 
boiling  the  solution,  a  current  of  carbon  dioxide  being  mean- 
while passed  into  it,  and  the  degree  of  reduction  was  determined 
by  titration  with  nearly  n/io  potassium  permanganate.  The 
solution  of  molybdic  acid  was  standardized  by  the  method  of 
Randall  J  and  also  by  evaporating  a  portion  to  dryness  and 

*  See  Randall,  page  426;  Gooch  and  Edgar,  page  349. 
'  t  Graham  Edgar,  Am.  Jour.  Sci.,  [4],  xxv,  332. 
J  See  page  426. 


428  METHODS  IN  CHEMICAL  ANALYSIS 

igniting  at  low  red  heat.  The  results  of  the  experiments  are 
given  in  the  table. 

The  results  show  that  if  the  concentration  be  not  greater  than 
0.2  grm.  of  MoO3  in  50  cm.3  of  solution,  and  the  acidity  not  less 
than  I  cm.3  of  sulphuric  acid  (sp.  gr.  1.84)  in  the  same  volume, 
the  molybdic  acid  is  not  reduced,  and  that  if  the  acidity  be 
increased  to  5  cm.3  of  sulphuric  acid,  reduction  does  not  occur  at 
even  a  concentration  of  0.4  grm.  in  25  cm.3. 

If  a  solution  containing  vanadic  acid  and  molybdic  acid  be 
treated  with  sulphur  dioxide  under  suitable  conditions,  only  the 
vanadic  acid  will  be  reduced 

V205  +  S02  =  V204  +  S03, 

and  the  degree  of  reduction,  determined  by  titration  with  per- 
manganate, will  measure  the  vanadic  acid,  according  to  the 
equation 

5  V204+2  KMn04+3  H2SO4  =  5  V2O5+K2SO4+MnSO4+3  H2O. 

If  the  solution  be  passed  through  the  zinc  column  both  the  va- 
nadic acid  and  the  molybdic  acid  will  be  reduced 


2  MoO3  +  3  Zn  =  Mo2O3  +  3  ZnO, 

and  oxidation  by  permanganate  will  take  place  according  to  the 
Aquations 


5  Mo2O3  +  6  KMnO4  +  9  H2SO4  =  10  MoO3  +  3  K2SO4 

+  6  MnSO4  +  9  H2O. 

If  the  number  of  centimeters  of  permanganate  used  in  the 
titration  of  the  vanadium  tetroxide  be  multiplied  by  three,  and 
the  product  subtracted  from  the  total  number  of  centimeters  of 
permanganate  used  in  the  titration  of  vanadium  dioxide  and 
molybdenum  sesquioxide,  the  result  is  the  number  of  centime- 
ters used  in  oxidizing  Mo2O3  to  MoO3,  from  which  the  amount 
of  molybdic  acid  present  may  be  easily  calculated. 

According  to  the  procedure  recommended  by  Edgar,  the  solu- 
tion containing  vanadic  and  molybdic  acids  is  diluted  to  75  cm.3, 
acidified  with  2.  cm.3  to  3  cm.3  of  strong  sulphuric  acid,  heated  to 
boiling  and  subjected  to  a  current  of  sulphur  dioxide  for  a  few 
minutes  until  the  clear  blue  color  indicates  the  complete  reduction 


MOLYBDENUM 


of  the  vanadic  acid  to  the  state  of  tetroxide.  The  boiling  is 
continued  for  some  time,  a  current  of  carbon  dioxide  being 
passed  into  the  liquid  until  the  last  trace  of  sulphur  dioxide  has 
been  removed.  Titration  is  then  effected  by  nearly  n/io  potas- 
sium permanganate  and  the  vanadic  acid  is  calculated. 

Next,  the  solution  just  titrated,  preceded  by  100  cm.3  of  hot 
water,  and  125  cm.3  of  dilute  sulphuric  acid  (2.5  per  cent),  and 
followed  by  100  cm.3  of  the  dilute  sulphuric  acid  and  then  by 
200  cm.3  of  hot  water,  is  passed  slowly  through  a  column  of 
amalgamated  zinc  in  a  Jones  reductor*  into  the  receiver  con- 
taining a  solution  of  ferric  alum,  and  the  hot  solution  is  titrated 
with  nearly  n/io  potassium  permanganate,  a  little  phosphoric 
acid  being  added  to  decolorize  the  ferric  salt.  The  amount  of 
permanganate  used  in  this  titration,  diminished  by  three  times 
the  amount  used  in  titrating  the  vanadium  tetroxide  reduced  by 
sulphur  dioxide,  measures  the  molybdic  acid. 

The  experimental  results  show  that  molybdic  acid  and  va- 
nadic acid  may  be  accurately  estimated  in  the  presence  of  one 
another  by  two  processes  of  reduction  and  oxidation,  the  reduc- 
tion being  made  first  by  sulphur  dioxide  and  last  by  amalga- 
mated zinc. 

Molybdic  and  Vanadic  Acids. 


KMnO4 
H/ioXi.o52. 

KMnO4 

M/IOXI.OS2 

VA 

taken  as 
NaVO3. 

MoO3 
taken  as 

VA 

found. 

Mo03 
found. 

Error  on 

VA. 

Error  on 
MoO3. 

cm.1 

cm.3 

grm. 

grm. 

grm. 

grm. 

grm. 

gnu. 

n-95* 

74-15 

0.1144 

0.1930 

0.1146 

0.1934 

+O  .  OOO2 

+0.0004 

n-95* 

74-00 

0.1144 

0.1930 

0.1146 

0.1926 

+O.OOO2 

—0.0004 

11.94* 

74.20 

0.1144 

0.1930 

0.1145 

0.1936 

+O.OOOI 

+o  .  0006 

5-97t 

37.10 

0.0572 

0.0965 

0.0572 

0.0965 

0.0000 

o  .  oooo 

5-97t 

37-05 

0.0572 

0.0965 

0.0572 

0.0962 

+0.0000 

—0.0003 

i?'98t 

37.12 

0.0572 

0.0965 

0-0573 

0.0963 

+O.OOOI 

—  O.OOO2 

55-0 

0.1144 

0.0965 

0.1146 

0.0967 

+O.OOO2 

+O.OOO2 

ii  .95! 

55-o 

0.1144 

0.0965 

0.1146 

0.0967 

+O.OOO2 

+O.OOO2 

n.M 

54.86 

0.1144 

0.0965 

0.1147 

0.0958 

+0.0003 

—  O.OOO7 

i7-92§ 

92.0 

0.1716 

0.1930 

0.1719 

0.1931 

+0.0003 

+O.OOOI 

i7-94§ 

92.05 

0.1716 

0.1930 

o.  1720 

0.1931 

+O.OOO4 

+O.OOOI 

17.9*1 

92.02 

0.1716 

0.1930 

0.1719 

0.1932 

+0.0003 

+0.0002: 

*  With  8  cm.»  of  sirupy  phosphoric  acid  and  50  cm.J  of  10  per  cent  ferric  alum, 
t  With  4  cm.1  of  sirupy  phosphoric  acid  and  25  cm.3  of  10  per  cent  ferric  alum. 
J  With  6  cm.*  of  sirupy  phosphoric  acid  and  35  cm.3  of  10  per  cent  ferric  alum. 
§  With  10  cm.*  of  sirupy  phosphoric  acid  and  65  cm.3  of  10  per  cent  ferric  alum. 

*  See  page  426. 


430  METHODS  IN  CHEMICAL  ANALYSIS 

URANIUM. 
The  Determination  of  Uranium  by  the  Aid  of  the  Jones  Reductor. 

In  studying  the  behavior  of  uranyl  sulphate  in  the  zinc*  reduc- 
tor  Pulman  *  has  shown  that  the  degree  of  reduction  varies  with 
conditions,  and  that  when  the  reduced  solution  is  received  in  an 
atmosphere  of  carbon  dioxide  the  titration  with  permanganate 
shows  that  the  uranium  salt  has  been  reduced  below  the  uranous 
stage,  represented  by  the  oxide  UO2,  but  that  the  over-reduction 
may  be  corrected  by  brief  contact  with  the  oxygen  of  the  air. 

The  details  of  procedure  are  as  follows:  The  uranium  sul- 
phate solution,  in  volume  100  cm.3  to  150  cm.3,  and  containing 
sulphuric  acid  in  the  proportion  [i  :  6],  is  heated  nearly  to  boil- 
ing. Preceded  by  a  few  cubic  centimeters  of  acid  of  the  same 
strength,  the  solution  is  drawn  by  gentle  suction  through  the 
1 8-inch  reductor  charged  with  20- mesh  amalgamated  zinc,  and  is 
followed  by  more  of  the  same  acid,  used  in  washing  the  container, 
and  250  cm.3  of  hot  water.  The  contents  of  the  receiving  flask 
are  poured  through  the  air  into  a  porcelain  dish,  diluted  with 
about  200  cm.3  of  hot  water,  and  titrated  with  n/io  potassium 
permanganate. 

The  proportion  of  free  sulphuric  acid  should  be  kept  during 
the  digestion  nearly  at  the  ratio  [i  :  6],  since  with  less  the  reduc- 
tion is  delayed,  while  more  produces  too  rapid  evolution  of 
hydrogen.  With  acid  in  the  ratio  [i  :  6],  fifteen  minutes  or  more 
should  be  taken  in  passing  uranium  sulphate  equivalent  to 
0.2  grm.  of  uranic  oxide  through  the  reductor;  for  0.3  grm.  half 
an  hour  or  more  should  be  allowed.  Care  is  taken  that  the 
liquid  in  the  reductor  shall  always  cover  the  zinc,  lest  hydrogen 
dioxide,  formed  by  contact  of  nascent  hydrogen  and  air,  vitiate 
the  results. 

The  contents  of  the  receiving  flask  after  the  reduction  are 
olive-green,  but  upon  exposure  to  the  air  by  pouring  into  the 
dish  the  color  changes  immediately  to  the  sea-green  color  always 
possessed  by  uranous  salts,  and  this  change  of  color  is  of  itself 
evidence  of  oxidation.  In  the  titration  of  the  hot  solution  of 
uranous  sulphate  with  permanganate  the  solution  gradually  be- 
comes more  and  more  yellowish  green  as  the  highest  condition 
of  oxidation  is  approached.  With  small  amounts  of  uranium 
*  O.  S.  Pulman,  Jr.,  Am.  Jour.  Sci.,  [4],  xvi,  229. 


URANIUM 


431 


the  addition  of  a  single  drop  of  permanganate  in  excess  brings 
out  a  faint  pink  color,  but  with  larger  amounts  the  end  point  is 
a  yellowish  pink* 

The  results  obtained  in  applying  this  method  to  uranyl  sul- 
phate prepared  from  the  pure  nitrate  by  evaporation  with  sul- 
phuric acid  are  shown  in  the  table. 

\ 

Reduction  in  the  Zinc  Column;  Exposure  to  Air;  Titration  with  Permanganate. 


Uranyl 

sulphate 
taken,  in 

H2S04 

Dilution 
at  time  of 

KMn04. 

Error  in 
terms  of 

terms  of 
U03. 

(1.84). 

reduction. 

Time. 

U03. 

grin. 

cm.3 

cm.3 

minutes. 

cm.3 

grm.  UO3. 

grm. 

0.1336 

18 

117 

15 

9-32 

0.1334 

—  O.OOO2 

0.1337 

20 

120 

15 

9-37 

0.1341 

+0.0004 

0.1336 

25 

125 

I? 

9.40 

O.I34S 

+o  .  0009 

0.2005 

18 

117 

2O 

14.02 

O.2Oo6 

+0.0001 

0.2003 

25 

125 

17 

14.01 

0.2005 

+O.OOO2 

0.2671 

23 

ISO 

2O 

18.67 

0.2671 

O.OOOO 

0.2673 

20 

140 

18 

18.65 

0.2669 

—  O.OOO4 

O.IOOI 

25 

I2S 

22 

7.06 

O.IOIO 

+O.OOO9 

O.  IOO2 

2O 

140 

17 

7.02 

0.1004 

+O.OOO2 

O.IOO2 

2O 

140 

14 

7.01 

o.  1003 

+O  .  OOOI 

0.0668 

18 

117 

17 

4-70 

0.0673 

+O.OOO5 

O.OQ94 

20 

IOO 

16 

6.96 

0.0995 

+O.OOOI 

0.1988 

20 

1  20 

18 

13.90 

0.1988 

o  .  oooo 

0.1988 

25 

150 

18 

13.88 

0.1985 

—0.0003 

0.33H 

25 

ISO 

27 

23-14 

0.3309 

—0.0005 

0.3314 

30 

150 

36 

23.19 

0.3316 

+O.OOO2 

0.3314 

25 

145 

33 

23.17 

0.3313 

—  O.OOOI 

CHAPTER  XL 


FLUORINE;   CHLORINE;  BROMINE;  IODINE. 

FLUORINE. 
The  Detection  of  Fluorine. 

Browning  *  has  shown  that  small  amounts  of  fluorine  may  be 
detected  by  the  converse  of  the  method  previously  described  for 
the  detection  of  silicates  and  fluosilicates.f  According  to  the 
procedure,  the  fluoride  is  put,  with  a  suitable  amount  of  silica, 
in  a  small  lead  cup,  i  cm.  in  diameter  and  depth  a  few  drops 
of  concentrated  sulphuric  acid  are  added ;  the  cup  is  covered  by 
a  flat  piece  of  lead  with  a  small  hole  in  the  center;  upon  the 
cover  is  placed  a  piece  of  moistened  black  filter  paper  and  upon 
this  a  small  pad  of  moistened  filter  paper  to  keep  the  black 
paper  moist  during  subsequent  heating  upon  the  steam  bath. 
After  about  ten  minutes'  heating  a  white  deposit  is  found  on  the 
under  side  of  the  black  paper,  over  the  opening  in  the  cover,  if 
fluorine  is  present  in  appreciable  amount. 

The  results  of  tests  are  given  in  the  table. 

Tests  for  Fluorine. 


Name  and  amount  of  fluoride  used, 
grm. 

Approximate 
per  cent  of  P. 

SiO2  present, 
grm. 

Result. 

o  1000  CaF2 

40 

None 

Nothing. 

o  oioo  CaF2 

40 

o  0500 

Very  good. 

o  0050  CaF2  

49 

0.0500 

Apparent. 

o.ooio  CaF2  

49 

0.0500 

Trace. 

o.oioo  NasAlFe  

54 

0.0500 

Very  good. 

o  0050  NaaAlFe 

C4 

o  0500 

Distinct. 

The  Acidimetric  Estimation  of  *Fluosilicic  Acid. 

Hileman  {  has  studied  methods  in  use  for  the  determination  of 
fluorine  in  fluosilicic  acid  by  neutralization  with  standard  alkali. 
The  first  set  of  these  methods  depends  upon  the  action  of  fluo- 

*  Philip  E.  Browning,  Am.  Jour.  Sci.,  [4],  xxxii.  249. 
t  See  page  241,  Fig.  20. 

J  Albert  Hileman,  Am.  Jour.  Sci.,  [4],  xxii,  329. 
432 


FLUORINE 


433 


silicic  acid  upon  potassium  chloride  in  alcoholic  solution  and 
titration  (without  removal  of  the  precipitated  potassium  fiuosil- 
icate)  of  the  liberated  hydrochloric  acid  by  standard  ammonia* 
or  by  standard  fixed  alkali,  f  The  second  set  of  methods  depends 
upon  the  titration  of  fluosilicic  acid  in  water  solution  by  standard 
alkali  hydroxide  to  the  complete  decomposition  of  the  fluosilicate. 
Hileman's  results  are  given  in  the  following  tables: 


Titration  in  Alcoholic  Solution 


2ROH  +  2HCI 


H,SiF,. 
cm.8 

Standard 
NH4OH. 

cm.8 

Standard 
KOH. 

cm.8 

Standard 
NaOH. 

cm.8 

Fluorine 
found. 

grm. 

Average, 
grm. 

25 

7-3 

0.1433 

25 

7-3 

0.1433 

25 

7.27 

0.1426 

0.1428 

25 

7-23 

0.1429 

25 

7-29 

O.I43I    • 

25 

10.67 

0.1412 

25 

10.72 

o.  1419 

25 

10.64 

0.1408 

O.I4II 

25 

10.67 

O.I4I2 

25 

10.60 

0.1403     , 

25 

. 



9.II 

0.1416    ] 

25 

.    . 

9.12 

O.I4I8 

25 

.    . 

9.07 

O.I4IO 

0.1415 

25 

.    . 

9.10 

0.1414 

25 

9.12 

0.1418    J 

The  differences  between  the  amounts  of  fluorine  indicated  by 
the  individual  determinations  in  any  one  of  these  processes  are 
generally  slight.  The  averages  of  the  determinations  by  potas- 
sium hydroxide  and  sodium  hydroxide  are  very  close  together, 
being  0.1411  grm.  and  0.1415  grm.  of  fluorine.  The  average  of 
the  titrations  by  ammonium  hydroxide  is  a  little  higher,  namely, 
0.1428  grm.  That  the  differences  between  these  averages  are 
due  to  gradual  variations  in  the  reading  tint  is  shown  by  a  com- 
parison of  three  titrations  as  nearly  simultaneous  as  possible,  in 
which  the  greatest  care  was  taken  to  bring  all  to  the  same  tint 
at  the  final  reading. 

*  Penfield,  Am.  Chem.  Jour.,  i,  27. 

t  Bullnheimer,  Zeit.  angew.  Chem.,  1901,  101. 


434  METHODS  IN  CHEMICAL  ANALYSIS 

Comparison  of  Simultaneous  Titrations  in  Alcoholic  Solution. 


Solution  used, 
cm.* 

Fluorine  found, 
grin. 

Titration  by  NH4OH 

7   2 

O   14.14. 

Titration  by  KOH  
Titration  by  NaOH  

10.71 
o   13 

o.  1418 

O   I4IQ 

\ 

It  appears  that  the  results  obtained  are  practically  the  same 
by  the  three  processes  of  neutralization  applied  to  a  solution  of 
fluosilicic  acid.  But  it  is  to  be  observed  that  all  are  possibly 
subject  to  a  common  and  constant  error  due  to  the  presence  of 
hydrofluoric  acid  as  well  as  fluosilicic  acid.  If  the  former  acid 
is  present  it  tends  to  raise  the  apparent  value  of  the  latter. 

With  these  results  of  titrations  in  alcoholic  solution  are  to  be 
compared  the  results  obtained  by  the  method  of  titration  in 
water  solution  (in  which  the  fluosilicate  is  completely  converted 
to  fluoride) ,  recorded  in  the  following  table  : 

Titration    of    Fluosilicic    Acid    in    Water    Solution: 
H2SiF6  +  6  ROH  =  6RF  +  SiO.H,  +  2  H20. 


H2SiF8  taken. 
cm.5 

Standard  KOH. 
cm.s 

Standard  NaOH. 
cm.' 

Fluorine  found, 
grm. 

Average, 
grm. 

25 
25 
25 

30.9 
30.8 

30.9 

0-I358 
O.I3S3 
0.1358 

1 

0.1355 

25 

30-79 

0.1353 

j 

25 



26.2 

0-1357 

1 

25 

26.15 

0.1355 

25 
25 

26.25 
26.2 

0.1360 
0.1357 

0.1358 

25 

26.13 

0.1354 

25 



26.14 

0.1354 

J 

It  is  obvious  that  the  process  of  titrating  fluosilicic  acid  in 
water  solution  yields  uniform  indications,  both  with  potassium 
hydroxide  and  sodium  hydroxide,  but  that  the  values  for  fluo- 
rine are  very  much  below  those  of  the  titrations  in  alcoholic 
solution.  And  this  will  be  the  case  if  the  solution  of  fluosilicic 
acid  contains  hydrofluoric  acid,  as  is  probable. 

The  action  of  ammonium  hydroxide  upon  fluosilicic  acid  in 
water  solution  proves  to  be  comparable  with  that  of  sodium 


FLUORINE 


435 


hydroxide,  and  inferentially  with  that  of  potassium  hydroxide, 
though  the  hydrolysis  of  the  fluosilicate  appears  to  be  not  quite 
so  complete. 

The  lodometric  Estimation  of  Fluosilicic  Acid. 

It  is  obvious  that  the  reaction  by  which  fluosilicic  acfd  liber- 
ates iodine  from  a  mixture  of  potassium  iodide  and  potassium 
iodate  may  be  turned  to  account  in  the  analysis  of  fluorides  as 
well  as  in  the  determination  of  fluosilicic  acid,*  provided  the 
course  of  action  is  regular. 

lodometric  and  Acidimetric  Determinations  in  Water  Solutions. 


H2SiF6. 

Standard  NaOH. 

Standard  Na2S2O3. 

Fluorine  found. 

Commercial  fluosilicic  acid. 


(i  cm.3  =0.005137  grm. 
of  fluorine.) 

(i  cm.3  =0.002335  grm. 
of  fluorine.) 

cm.3 

cm.3 

cm.3  . 

grm. 

2< 

10.82 

O.O^6 

25 

10.87 

0-0559 

25 

10.83 

0-0557 

2C 

27    C4 

O   CX4.Q 

25 



23-52 

0.0549 

25 

23.48 

0.0548 

25 

23-50 

o  .  0548 

25 

23-45 

o  .  0548 

25 

23.48 

0.0548 

Fluosilicic  acid  made  by  HF  on  excess  of 


(i  cm.3=o.oosi37  grm. 
of  fluorine.) 

(i  cm.3=o.oo244i  grm. 
of  fluorine.) 

25 
25 
25 

2< 

9.08 
9-05 
9-05 

18  78 

o  .  0466 
o  .  0464 
o  .  0464 

o  04.18 

25 



18.80 

0.0459 

Upon  testing  the  action  of  fluosilicic  acid  upon  the  iodide- 
iodate  mixture,  Hilemanf  has  found  that,  while  iodine  is  liberated 

*  See  page  436. 

t  Albert  Hileman,  Am.  Jour.  Sci.,  [4],  xxii,  383. 


436  METHODS  IN  CHEMICAL  ANALYSIS 

freely  in  the  cold,  a  complete  reaction  is  not  obtained  in  the 
course  of  several  hours  —  the  amount  of  iodine  liberated  indi- 
cating that  an  acid  other  than  fluosilicic  acid  is  acting.  It 
appears,  further,  that  on  heating  the  mixture  to  the  boiling 
point  in  a  flask  closed  with  a  glass  stopper  and  trapped  with  a 
solution  of  potassium  iodide,  nearly  one  equivalent  of  iodine  is 
liberated  for  every  equivalent  of  fluorine  present  as  fluosilicic 
acid,  according  to  the  reaction 

5  KI  +  KI03  +  H2SiF6  =  6  KF  +  3  I2  +  SiO2  +  H2O. 


The  Estimation  of  Fluorine  Evolved  as  Silicon  Fluoride. 

As  to  sources  of  error  in  the  determination  of  fluorine  by  the 
silicon  fluoride  processes,  due  to  imperfect  elimination  and  col- 
lection of  silicon  fluoride,  there  is  the  testimony  of  many  investi- 
gators. The  importance  of  using  the  fluoride  in  the  finest  state 
of  division,  of  having  the  sulphuric  acid  of  highest  strength,  of 
properly  absorbing  the  vapors  of  sulphuric  acid  evolved  from 
the  decomposition  flask,  and  of  using  quartz  for  the  silicon  dioxide 
in  the  decomposition  flask,  have  all  been  emphasized.  Many 
forms  of  apparatus  have  been  employed  and  the  results  have 
varied  widely. 
Elimination  of  In  consequence  of  difficulty  with  the  silicon 

Silicon  Fluoride    n          -j  •  1*1-11  ••• 

-      fluoride   processes   in   which   the  decomposition   is 


peratures.  effected  at  the  usual  temperatures,  between  150°  and 
1  60°,  Hileman  *  has  devised  a  convenient  form  of  apparatus  in 
which  the  acid  mixture  may  be  heated  to  boiling  to  facilitate  the 
removal  of  the  silicon  fluoride  to  the  absorption  system. 

A  glass  stopper,  made  by  drawing  out  a  glass  tube  I  cm.  in 
diameter  and  sealing  a  small  glass  tube  on  each  end,  is  ground 
into  a  70  cm.3  side-neck  flask.  To  one  end  is  sealed  a  glass  stop- 
cock. The  other  end  extends  to  the  bottom  of  the  flask.  The 
side  neck  is  sealed  to  a  Voit  flask.  The  length  of  the  tube 
between  the  two  flasks  is  17  cm.,  and  it  is«bent  at  a  point  about 
12  cm.  from  the  Voit  flask.  The  tube  leading  from  the  Voit 
flask  is  joined  to  a  large  empty  U-tube  and  this  is  connected, 
through  a  tube  charged  with  phosphorus  pentoxide,  with  the 
absorption  tube.  The  absorption  apparatus  is  similar  to  that 

*  Albert  Hileman,  Am.  Jour.  Sci.,  [4],  xxii,  329. 


FLUORINE 


437 


described  by  Burk,*  and  consists  of  a  test  tube  34  cm.  in  length 
and  2  cm.  in  diameter,  conveniently  inclined  and  containing  a 
few  cubic  centimeters  of  mercury  into  which  extends  an  inlet  tube 
with  a  capillary  opening  through  which  the  gas  is  delivered  under 
the  mercury  and  allowed  to  bubble  into  the  water  with  which 
the  tube  is  charged.  The  tube  is  connected  with  a  suction  pump 
by  means  of  a  T-tube  joined  also  to  an  air- trap  of  mercury  to 
regulate  the  pressure  within  the  apparatus  to  any  desired  degree. 


Fig.  28. 

Preparatory  to  making  a  determination  the  apparatus  is  care- 
fully dried.  The  tip  of  the  inlet  tube  is  then  placed  beneath 
the  surface  of  mercury  in  the  bottom  of  the  absorption  tube  and 
distilled  water  is  added,  care  being  taken  that  enough  space  shall 
remain  to  allow  for  the  rise  in  level  when  air  bubbles  through 
the  liquid.  The  pressure  regulator  is  adjusted  so  that  the  suc- 
tion shall  produce  slightly  diminished  pressure  in  the  apparatus. 
The  U-tube  is  immersed  in  a  vessel  of  cold  water  and  connected 
with  the  system. 

Next,  the  mineral,  together  with  quartz  powder  to  about  three 
times  the  weight  of  the  fluorine  present,  is  transferred  to  the 
decomposition  flask,  and  enough  sulphuric  acid  to  seal  the 
delivery  tube  from  the  side-neck  flask  is  introduced  into  the  Voit 
flask.  The  two  flasks  are  then  tilted  so  that  the  acid  shall 
moisten  the  connecting  tube  to  the  bend.  About  40  cm.3  of 
sulphuric  acid,  previously  boiled  for  half  an  hour  and  cooled  in  a 
current  of  dry  air,  and  several  capillary  tubes  of  the  form  recom- 
mended by  Scudderf  to  prevent  bumping,  are  put  in  the  decom- 

*  Jour.  Am.  Chem.  Soc.,  xxiii,  825. 
t  Ibid.,  xxv,  113. 


438 


METHODS   IN   CHEMICAL  ANALYSIS 


position  flask.  The  stopper  is  quickly  replaced  and  sealed  with 
a  drop  of  sulphuric  acid.  A  thin  strip  of  asbestos  is  wrapped 
about  the  neck  of  the  flask  and,  the  stopcock  having  been  closed, 
the  bulb  is  heated  in  a  radiator  covered  over  with  sheet  asbestos. 
When  the  heat  is  applied,  bubbles  of  gas  are  given  off,  the 
solid  material  rises  to  the  surface,  and  during  the  course  of  the 
heating  an  oily  film  gathers  on  the  upper  part  of  the  flask  and  in 
the  delivery  tube.  On  boiling,  this  film  is  replaced  by  a  white 
deposit  which  recedes  before  the  acid  vapors.  The  success  of 
the  determination  depends,  as  was  found,  on  the  breaking  up  of 
this  deposit,  which  is  probably  a  product  of  the  hydrolysis  of 
silicon  fluoride.  On  this  account  the  tube  between  the  two  flasks 
should  be  as  short  as  is  practicable.  Unsatisfactory  results  were 
obtained  when  the  tube  was  about  one-half  longer  than  the  di- 
mensions given  above.  When  the  acid  vapors  have  penetrated 
the  length  of  the  tube,  leaving  it  clear  or  translucent,  the  decom- 
position is  complete  and,  the  stopcock  having  been  opened,  the 
side-neck  flask  is  cooled  to  about  75°. 

Elimination  of  Silicon  Fluoride;  Absorption  in  Water;   Titration  of  Fluosilicic 
Acid  with  Sodium  Hydroxide. 


CaF2. 
grm. 

Quartz, 
grm. 

NaOH. 
cm.3 

Theory, 
fluorine. 

grm. 

Found, 
fluorine. 

grm. 

Error, 
fluorine. 

grm. 

0.3000 

0.4 

28.2 

0.1459 

0.1444 

—  0.0015 

o  .  3000 

0.4 

28.35 

Q-I459 

0.1452 

—  0.0007 

0.3000 

0.4 

28.34 

0.1459 

0.1451 

—  0.0008 

0.3000 

0.4 

23-5 

0.1215 

o  .  i  203 

—  O.OOI2 

0.3000 

0.4 

28.3 

o.i459 

o.  1419 

—  O.OOIO 

0.3000 

0.4 

28.37 

Q.I459 

0.1453 

—  0.0006 

NaF. 

0.3000 

0.4 

26.2 

0.1356 

0.1342 

—  O.OOI2 

0.3000 

0.4 

26.44 

0.1356 

0.1355 

—  o.oooi 

0.3000 

0.4 

26.46 

0.1356 

0.1356 

0.0000 

0.3000 

0.4 

26.3 

0.1356 

0.1347 

—0.0009 

0.3000 

0.4 

26.34 

0.1356 

0.1349 

—0.0007 

0.3000 

0.4 

26.3 

0.1356 

0.1347 

—0.0009 

Ignited 

Silicic  acid 

0.3000 

0.4 

26.29 

0.1356 

0.1346 

—  O.OOIO 

Quartz 

0.3000 

0.4 

26.3 

0.1356 

0,  1347 

—  0.0009 

If  the  acid  tends  to  suck  back  from  the  Voit  flask,  it  is  arrested 
by  opening  the  stopcock  to  relieve  the  vacuum,  and  it  is  at  this 
point  of  the  experiment  that  the  necessity  for  the  previous  ad- 


FLUORINE  439 

justment  of  the  pressure  regulator,  to  avoid  sudden  transfer  of 
acid  vapor  to  the  absorption  tube,  becomes  apparent.  The  de- 
composition is  ended  in  from  fifteen  to  forty  minutes.  A  current 
of  purified  air  is  drawn  through  the  apparatus,  slowly  at  first 
and  then  more  rapidly.  About  six  liters  are  necessary  to  re- 
move the  last  trace  of  silicon  fluoride.  The  delivery  tube  is 
washed  and  the  absorption  solution  transferred  to  a  flask  and 
titrated  with  sodium  hydroxide  prepared  from  sodium*,  with 
phenolphthalein  as  an  indicator,  according  to  the  following 
reaction :  f 

H2SiF6  +  6  NaOH  =  6  NaF  +  Si(OH)4  +  2  H2O. 

The  results  obtained  are  shown  in  the  preceding  table. 

With  the  apparatus  described  above,  in  which  the  sulphuric 
acid  in  the  decomposition  flask  may  be  boiled,  the  silicon  fluoride 
formed  passes  rapidly  to  the  absorption  system,  other  products 
of  partial  hydrolysis  of  silicon  fluoride  formed  in  the  flask  or 
tube  are  ultimately  reconverted  to  silicon  fluoride,  and  regular 
results  of  a  fair  degree  of  accuracy  are  obtained.  In  all  the 
experiments  except  two,  phosphorus  pentoxide  (about  2.0  grm.) 
was  introduced  into  the  decomposition  flask,  the  purpose  being 
to  retain  water  formed  in  the  reaction.  The  results  of  these  two 
experiments  show,  however,  that  phosphorus  pentoxide  in  the 
flask  is  not  essential. 

In  three  blank  experiments,  in  which  the  acid  in  the  decompo- 
sition flask  was  heated  to  boiling,  amounts  of  acid  in  each  case 
equivalent  to  0.0002  grm.  of  fluorine  were  found  in  the  absorp- 
tion solution;  the  results,  therefore,  are  subject  to  this  trifling 
error. 

lodometric  De-  Hileman  J  has  also  applied  the  iodometric  deter- 
Fiuari^iT  °f  mination  of  fluosilicic  acid  §  to  the  estimation  of 
Fluorides.  fluorine  in  silicon  fluoride  evolved  and  absorbed  in 
the  manner  described  above. 

Silicon  fluoride,  evolved  from  calcium  fluoride  in  the  apparatus 
previously  figured  and  by  the  method  described, ||  is  absorbed 
in  water,  and,  after  separating  the  mercury  used  in  the  apparatus 

*  Kiister,  Zeit.  anorg.  Chem.,  xli,  475. 

t  See  page  434. 

J  Am.  Jour.  Sci.,  [4],  xxii,  383. 

§  See  page  435. 

||  See  page  436. 


440 


METHODS  IN  CHEMICAL  ANALYSIS 


by  means  of  a  separating  funnel,  the  iodide-iodate  mixture  is 
added  to  the  solution  and  the  iodine  liberated  is  titrated  by 
sodium  thiosulphate.  The  results  are  given  below : 

Elimination  of  Silicon  Fluoride;  Absorption  in  Water;   lodometric  Estimation 

of  Fluosilicic  Acid. 


CaF,. 

Na2S203. 

Theory, 
fluorine. 

Found, 
fluorine. 

Error, 
fluorine. 

grm. 

cm.3 

grm. 

grm. 

grm. 

0.2500 
0.2300 

51-35 
46.75 

0.1216 
0.1119 

•    0.1199 
0.1091 

0.0017 
O.OO28 

0.2300 

NaF. 

47-15 

0.1119 

O.  IIOO 

0.0019 

O.2OOO 

38.50 

0.0903 

o  .  0899 

O.OOI4 

The  average  error  of    —0.0019  grm.  is  considerably  greater 
than  that  of  the  acidimetric  method,  —0.0008  grm. 


CHLORINE;   BROMINE;  IODINE. 

The  Detection  of  Iodine,  Bromine  and  Chlorine  in  Presence  of  One 

Another. 

In  the  qualitative  testing  of  substances  it  is  not  a  matter  of 
moment  that  a  portion  of  a  substance  looked  for  escapes  the  re- 
action, provided  enough  is  left  to  furnish  the  indication  sought. 
In  the  separation  of  iodine  from  bromine  and  chlorine,  Gooch 
and  Brooks  *  have  applied  the  nitrite  reaction  f  to  small  amounts 
of  liquid  in  test  tubes  and  have  demonstrated  that  the  losses  of 
chlorine  and  bromine  under  the  conditions  are  proportioned  to 
the  strength  of  the  solution;  or,  in  other  words,  that  when  the 
amounts  of  bromine  and  chlorine  are  very  small  they  escape 
volatilization,  and  when  large  a  sufficient  amount  remains  to  give 
strong  tests. 

The  process  elaborated  for  the  rapid  qualitative  detection  of 
the  halogens  in  presence  of  one  another  may  be  summarized  as 
follows : 

To  detect  iodine,  the  solution  of  the  substance  under  exami- 
nation is  acidulated  with  dilute  sulphuric  acid  and  treated  with 
a  drop  or  two  of  a  solution  of  sodium  or  potassium  nitrite  free 

*  F.  A.  Gooch  and  F.  T.  Brooks,  Am.  Jour.  Sci.,  fc],  xl,  283. 
t  See  page  451- 


CHLORINE;   BROMINE;  IODINE  441 

from  chlorine.  Unless  the  amount  present  is  small,  the  iodine 
shows  itself  in  the  color  of  the  solution  and  in  the  vapors  which 
escape.  Small  amounts  may  be  found  by  shaking  the  liquid 
with  carbon  disulphide  in  the  usual  manner,  or,  when  economy 
of  material  is  desirable,  by  gently  heating  the  prepared  solution 
and  testing  the  escaping  fumes  with  red  litmus  paper,  thus 
utilizing  the  same  portion  of  material  for  the  detection  of  the 
iodine  and  for  its  separation  preparatory  to  testing  for  bromine 
and  chlorine. 

To  remove  the  iodine  previous  to  making  the  tests  for  bromine 
and  chlorine,  a  few  drops  of  dilute  sulphuric  acid  and  a  like  amount 
of  a  dilute  solution  of  sodium  or  potassium  nitrite  (prepared 
free  from  chlorine  by  adding  a  little  silver  nitrate,  faintly  acidu- 
lating with  nitric  acid,  and  filtering)  are  added  to  the  solution  of 
the  substance  in  a  test  tube,  and  the  liquid  is  boiled  with  constant 
agitation.  When  the  color  of  iodine  disappears  from  the  fumes 
and  the  solution,  a  drop  or  two  of  sulphuric  acid,  and  of  the 
nitrite,  are  again  added,  and  the  boiling  is  repeated.  When 
the  escaping  steam  no  longer  gives  to  red  litmus  paper  the  char- 
acteristic gray  blue  color  due  to  the  action  of  iodine,  the  process 
of  separation  is  complete. 

A  portion  of  the  solution  thus  prepared  is  tested  for  bromine 
by  cautiously  adding  a  dilute  solution  of  sodium  hypochlorite 
and  shaking  with  colorless  carbon  disulphide. 

The  test  for  chlorine  is  made  in  a  second  portion  of  the  solu- 
tion from  which  the  iodine,  has  been  removed.  The  liquid  is 
neutralized  with  sodium  carbonate  or  hydroxide  free  from  chlo- 
ride, evaporated  to  dry  ness  in  a  test  tube  and  treated  with 
sulphuric  acid  and  potassium  dichromate,  the  fumes 
of  the  chlorochromic  anhydride  which  arise  on  gentle 
warming  being  condensed  and  converted  to  chromic 
acid  by  a  film  of  moisture  upon  the  interior  walls  of 
a  trap,  such  as  is  shown  in  Fig.  6,*  or  in  a  mushroom 
trap,  shown  in  Fig.  29.  The  trap  is  washed  out  with 
a  very  little  distilled  water  (5  cm.3)  and  the  washjngs, 
made  slightly  ammoniacal  to  destroy  free  bromine,  if 
necessary,  and  after  gentle  warming  again  acidified, 
are  tested  with  lead  acetate.  If  yellow  lead  chromate  is  pre- 
cipitated the  presence  of  chlorine  in  the  original  substance  is 

*  See  page  6. 


442 


METHODS  IN  CHEMICAL  ANALYSIS 


proved.  If  the  precipitate  is  white,  as  is  very  likely  to  be  the 
case,  a  few  drops  of  a  saturated  solution  of  ammonium  acetate  are 
added  with  caution,  and  the  whole  is  gently  warmed  to  dissolve 
the  white  sulphate.  On  cooling  the  solution,  and  shaking  (or 
immediately,  if  much  chromic  acid  has  been  formed),  the  yellow 
chromate  falls,  or  gives  color  to  the  solution,  according  as  the 
chloride  was  originally  present  in  large  or  small  amount. 

The  process  is  rapid  and  sufficiently  exact  for  qualitative  test- 
ing in  general.  The  results  of  experiments  made  to  determine 
the  delicacy  of  the  tests  for  bromine  and  chlorine  in  a  substance 
treated  to  remove  iodine  are  shown  in  the  tabular  statements : 


Hypochlorite  Test  for  Bromine  after  Treatment  to  Remove  Iodine. 


KI  taken, 
gnu. 

KBr  taken, 
grm. 

Total 
volume. 

cm.8 

Ratio  of  KBr  to 
solution. 

Color  in  carbon 
disulphide. 

O.IOOO 

0.0010 

10 

10,000 

Pronounced. 

0.1000 

0.0004 

IO 

25,000 

Pronounced. 

O.IOOO 

0.0003 

IO 

33,000 

Pronounced. 

O.IOOO 

O.OOO2 

IO 

50,000 

Faint. 

O.IOOO 

O.OOO2 

IO 

50,000 

Faint. 

O.IOOO 

0.0001 

1°  . 

100,000 

None. 

0.0500 

0.0005 

IO 

20,000 

Pronounced. 

0.0400 

0.0004 

IO 

25,000 

Pronounced. 

0.0300 

0.0003 

IO 

33-000 

Pronounced. 

O.O2OO 

O.OOO2 

10 

50,000 

Faint. 

O.OIOO 

0.0001 

10 

100,000 

None. 

O.IOOO 

0.00007 

5 

70,000 

.  Trace. 

0.0070 

0.00007 

5 

70,000 

Trace. 

Chromate  Test  for  Chlorine  after  Treatment  to  Remove  Iodine. 


KI  taken, 
grm. 

KBr  taken. 

grm. 

KC1  taken, 
grm. 

Final  volume. 
cm.8 

Reaction  obtained. 

0. 

0.0030 

5 

Marked  precipitation. 

. 

o. 

O.OO2O 

5 

Distinct  precipitation. 

. 

O. 

O.OOIO 

5 

Distinct  color. 

o. 

0.0005 

5 

Faint  color. 

o. 

0.0005 

5 

Faint  color. 

0. 

O.OOO4 

5 

Faintest  color. 

o. 

0.0003 

5 

Doubtful  color. 

o. 

0.0002 

5 

Doubtful  color. 

o. 

O.OOOI 

5 

None. 

O.I 

0. 

O.OOIO 

5 

Distinct  color. 

O.  I 

o. 

O.OOIO 

5 

Distinct  color. 

O.I 

o. 

O.OOIO 

5 

Distinct  color. 

CHLORINE;   BROMINE;   IODINE 


443 


The  evolution  of  considerable  amounts  of  bromine  appears  to 
diminish  the  delicacy  of  the  test  in  some  degree,  but  0.0005  grm- 
of  chlorine  —  the  amount  in  o.ooio  grm.  of  potassium  chloride  — 
is  indicated  unmistakably  in  the  presence  of  o.i  grm.  of  potas- 
sium bromide,  and  o.i  grm.  of  potassium  iodide,  and  the  test 
may  probably  be  relied  upon  to  show  half  that  amount  of 
chlorine. 


The  Determination  of  Free  Chlorine  and  Free  Bromine  by  Liber- 
ation of  Iodine  and  Absorption  of  that  Element  by  Silver. 

Chlorine  and  bromine  in  free  condition  may  be  determined  by 
allowing  them  to  act  upon  potassium  iodide  in  solution  made 
acid  with  hydrochloric  acid  and  in  an  atmosphere  of  hydrogen, 
shaking  the  mixture  with  a  weighed  amount  of  specially  prepared 
silver,  and  determining  the  increase  in  weight  of  the  silver  due 
to  absorption  of  iodine  set  free,  according  to  the  method  de- 
scribed for  the  determination  of  iodine  by  absorption  in  metallic 
silver.*  Tests  of  this  process  by  Perkins  f  were  made  upon 
definite  amounts  of  aqueous  solutions  of  chlorine  and  bromine 
determined  titrimetrically  at  the  time  of  the  experiments.  These 
definite  amounts  of  the  aqueous  solutions  were  put  in  a  flask 
containing  an  excess  of  potassium  iodide  made  acid  with  hydro- 

Liberation  of  Iodine  and  Absorption  by  Silver. 


Silver  taken, 
grm. 

Halogen  taken, 
grm. 

Iodine  found, 
grm. 

Calculated  amount 
of  halogen. 

grm. 

Error, 
grm. 

Determination  of  bromine. 


3.0000 

0.0213 

0.0336 

O.O2II 

—  O.OOO2 

3  .  oooo 

0.0426 

0.0678 

0.0427 

-f-o.oooi 

3.0000 

0.1065 

0.1694 

o.  1067 

+O.OOO2 

Determination  of  chlorine. 


3.0000 

0.0161 

0-0574 

0.0160 

—  O.OOOI 

3.0000 

0.0322 

o.  1146 

0.0320 

—  0.0002 

3.0000 

0.0322 

0.1145 

0.0320 

—  O.OOO2 

3.0000 

0.0322 

0.1141 

0.0318 

—  0.0004 

3.0000 

0.0483 

o.  1716 

o  .  0479 

—  O.OOO4 

*  See  page  444. 

t  Claude  C.  Perkins,  Am.  Jour.  Sci.,  [4],  xxix,  338. 


444  METHODS  IN  CHEMICAL  ANALYSIS 

chloric  acid  in  an  atmosphere  of  hydrogen,  and  the  mixture 
was  shaken  with  a  weighed  amount  of  the  specially  pre- 
pared silver.  The  residue  of  silver  and  silver  iodide  was  col- 
lected on  asbestos  in  a  perforated  crucible,  washed,  dried  and 
weighed.  The  increase  in  weight  represents  the  weight  of 
iodine  liberated  and  from  this  the  amount  of  chlorine  or  bro- 
mine is  calculated.  Results  of  this  process  are  given  in  the 
table. 


The  Gravimetric  Determination  of  Iodine  by  Absorption  in  Metallic 

Silver. 

It  has  been  shown  by  Gooch  and  Perkins  *  that 

Free  Iodine.  .  J 

tree  iodine  may  be  determined  with  accuracy  in 
solution  in  potassium  iodide,  either  neutral  or  alkaline  with  an 
acid  carbonate,  by  shaking  the  solution  with  metallic  silver  in  a 
closed  flask  rilled  with  hydrogen  and  determining  the  increase 
in  weight  of  the  silver.  Silver  reduced  from  a  silver  salt  by  zinc, 
or  from  silver  sulphide  by  hydrogen,  may  serve  the  purpose,  pro- 
vided it  is  subjected  to  a  preliminary  treatment  with  potassium 
iodide,  and  silver  reduced  from  the  oxide  by  hydrogen  is  also 
serviceable;  but  the  best  form  of  silver,  and  the  one  most  easily 
prepared  in  the  pure  state,  is  that  deposited  electrolytically  upon 
a  small  oscillating  cathode  of  platinum  from  a  solution  of  silver 
nitrate,  the  platinum  anode  being  inclosed  in  a  porous  cell. 
The  shaking  of  the  silver  may  be  done  by  hand  or  by  some 
simple  form  of  mechanical  shaker  like  that  described  and  figured 
elsewhere. f  In  the  test  experiments,  detailed  in  the  table,  n/io 
iodine  solution  was  drawn  from  a  burette  into  a  250  cm.3  Erlen- 
meyer  flask  containing  a  weighed  amount  of  finely  divided  silver. 
The  flask,  properly  trapped  and  attached  to  a  mechanical  shaker 
adjusted  to  give  the  liquid  a  rapid  rotary  motion,  was  shaken 
until  the  iodine  color  had  vanished.  The  liquid,  usually  50  cm.3 
in  volume,  was  diluted  to  about  100  cm.3,  and  the  residue  of  silver 
and  silver  iodide,  collected  in  a  perforated  crucible  fitted  with 
asbestos,  was  washed,  dried  between  130°  and  140°,  and  weighed. 
The  difference  between  the  weight  of  silver  taken  and  that  of 
the  residue  of  silver  and  silver  iodide  should,  according  to  the 

*  F.  A.  Gooch  and  Claude  C.  Perkins,  Am.  Jour.  Sci.,  [4],  xxviii,  33. 
t  See  page  9. 


CHLORINE;   BROMINE;  IODINE 


445 


The  Action  of  Silver  upon  n/io  Iodine  in  an  Atmosphere  of  Hydrogen. 


Silver  taken, 
grm. 

Iodine  taken, 
grm. 

Increase  in  weight 
of  iodine. 

grm. 

Error  in  iodine, 
grm. 

Average  error  in 
iodine. 

grm* 

A.   Silver  reduced  from  AgCl  by  zinc  and  treated  with  KI. 


3.0000 
i  .0000 

0.6461 
0-6447 

o  .  6464 
o  .  6448 

+0.0003 

+O.OOOI 

+O.OOO2 

B.   Silver  reduced  from  Agl  by  zinc  and  treated  with  KI. 

3.6293 
3  .  2049 
3.0000 
3.0068 
3.0049 
3.0026 
2.9990 
3.0005 

0.3217 
0.3217 
0.3217 
0.3217 
0.3217 
0-6434 
0.3217 
0.3217 

0.3221 
0.3225 
0.3219 
0.3212 
0.3221 

0.6441 

0.3214 
0.3214 

+0  .  0004 
-j-o  .  0008 
+0.0002 
—0.0005 
+o  .  0004 
+0.0007 
—0.0003 
—0.0003 

+O.OOO2 

%C.   Silver  reduced  from  Ag2S  by  hydrogen  and  treated  with  KI. 

3.0000 
3.0000 

0.6461 
0.6461 

0.6463 

o  .  6460 

+0.0002 

—  o.oooi 

+0.0001 

D.   Silver  reduced  from  Ag2O  by  hydrogen. 

3.0000 
3.0000 

0.6434 
0-6434 

0.6443 

o  .  6430 

+0.0009 
—0.0004 

+o  .  0003 

E.   Silver  reduced  electrolytically  from  AgNO3upon  an  oscillating  cathode. 

4.4189 
3.0025 
3.0009 

3-oi57 
3-0000 
3.0000 
3.0004 
3-0043 
3.0000 
3-58io 
3.0000 

0.6447 
0-6447 
0-6447 
0.6447 
0-6447 
0.6447 
0-6447 
0.6447 
o  -  6434 
0.3217 
0.3217 

0.6447 

o  .  6448 

0.6443 
0.6445 
o  .  6444 
0.6452 
0.6443 
0.6443 
0.6430 
0.3221 
0.3219 

o.oooo 

+O.OOOI 

—0.0004 

—  O.OOO2 

—0.0003 
+o  .  0005 
—0.0004 
—0.0004 
—0.0004 
+o  .  0004 

+O.OOO2 

—O.OOOI 

Electrolytic  silver  in  presence  of  NaHCO3. 

3.0014 
3.0169 
3-0083 
3.0016 
3.0069 

0.3217 
0.3217 
0.6434 
0.2500 
0.3217 

0.3216 
0.3216 
0.6433 
0.2503 
0.3219 

—  O.OOOI 
—  O.OOOI 

—o.oooi 
+0.0003 

+O.OOO2 

+O.OOOI 

446 


METHODS  IN  CHEMICAL  ANALYSIS 


theory  of  action,  be  the  measure  of  the  free  iodine.  The  time 
required  for  the  absorption  of  approximately  0.65  grm.  of  iodine 
in  50  cm.3  of  liquid  was  from  fifteen  to  twenty-five  minutes.  The 
mean  error  of  sixteen  determinations  in  which  electrolytic  silver 
was  employed  proved  to  be  0.00004  grm-  between  extremes  of 
-1-0.0005  grm.  and  —0.0004  grm- 

[  This  process,  in  which  free  iodine  is  absorbed  by  specially 
prepared  silver*  under  hydrogen,  either  in  neutral  solution 
or  in  a  solution  made  alkaline  with  an  acid  carbonate,  is 
applicable  in  many  analytical  operations  involving  liberation 
of  iodine  as  well  as  in  the  gravimetric  standardization  of  the 
usual  iodine  solution  of  volumetric  analysis.  The  adaptation 
of  the  process  to  the  determination  of.  iodine  in  iodides,  free 
chlorine,  free  bromine,  and  various  oxidizers,  has  been  described 
by  Per  kins,  f 

In  applying  the  process  to  the  determination  of 
combined  iodine  a  definite  amount  of  the  iodide  in 
solution  is  introduced  into  a  flask  with  an  excess  over  the  calcu- 
lated amount  of  an  oxidizing  reagent  (usually  potassium  nitrite 
or  hydrogen  peroxide),  and  the  whole,  made  acid  with  hydro- 
chloric acid,  is  shaken  with  a  weighed  amount  of  silver.  The 
increase  in  the  weight  of  the  silver  indicates  the  amount  of 
Iodine  liberated,  and  from  this  the  amount  of  potassium  iodide 
may  be  easily  calculated.  The  table  shows  the  results  of  deter- 
minations with  potassium  iodide  in  solution  previously  stand- 
ardized by  the  distillation  method  with  sulphuric  acid  and 
potassium  hydrogen  arsenate.f 


Iodine  in  Iodides. 


Determination  of  the  Iodine  of  Potassium  Iodide. 


Silver  taken. 

KI  taken. 

Iodine  found. 

Calculated  amount 
of  KI. 

Error. 

grm. 

grm. 

grm. 

grm. 

grm. 

2.7803 

0.1144 

0.0872 

o.  1141 

—  0.0003 

3.0028 

0.1346 

0.1026 

0.1342 

—  0.0004 

2  .  7800 

0.1279 

0.0978 

0.1281 

4-O.OOO2 

2  .  0008 

0.1279 

0-0974 

0.1274 

—  0.0005 

3-0001 

0.1346 

0.1029 

0.1346 

O.OOOO 

*  See  page  27. 

f  See  also  pages  443,  444,  361. 

t  See  page  457. 


CHLORINE;   BROMINE;  IODINE  447 

The  Determination  of  Halogens  in  Benzol  Derivatives  by  the  Use 
of  Metallic  Potassium. 

The  method  of  Stephanoff  *  for  the  estimation  of  halogens  in 
aromatic  compounds,  by  reduction  with  sodium  and  alcohol,  was 
found  by  Maryott  f  to  give  irregular  and  low  results  when  ap- 
plied to  chlorbenzol.  The  action  of  the  sodium  is  very  slow, 
especially  after  the  liquid  has  become  pasty  owing  to  separation 
of  sodium  ethylate,  and  this  difficulty  is  only  partly  overcome  by 
increasing  the  amount  of  alcohol  on  account  of  the  unfavorable 
effect  of  dilution. 

Maryott  has  shown,  however,  that  a  similar  method,  based  on 
the  use  of  the  more  active  potassium,  in  place  of  sodium,  as  the 
reducing  agent,  gives  a  complete  reduction  and  accurate  ana- 
lytical results  when  employed  upon  the  halogen  substituted  benzols. 

As  the  action  of  potassium  upon  98  per  cent  alcohol  is  very 
energetic,  the  alcohol  used  is  diluted  with  twice  its  own  volume 
of  thiophene-free  benzol.  Very  little  heating  is  required,  so  that 
a  plain  glass  tube,  about  50  cm.  in  length,  serves  as  a  return 
cooler  in  place  of  the  water  cooled  condenser  used  by  Stephanoff. 

The  substance  to  be  analyzed  is  weighed  out  in  an  Erlenmeyer 
flask,  10  cm.3  to  15  cm.3  of  the  1:2  alcohol-benzol  mixture  are 
added,  the  cooling  tube  is  attached,  and  the  potassium,  in  small 
pieces,  is  gradually  dropped  in  through  the  tube.  The  weight 
of  potassium  required  is  about  ten  times  that  called  for  by  the 
equation 

C6H6C1-+  2  K  +  C2H5OH  =  C6H6  +  KC1  +  C2H5OK. 

The  greater  efficiency  of  potassium  as  a  reducer  as  compared 
with  sodium  is  shown  by  the  fact  that  an  analysis  of  chlorben- 
zol, carried  out  in  exactly  the  same  way  as  the  analyses  tabu- 
lated below  except  that  sodium  (ten  times  the  theoretical 
amount)  was  used  instead  of  potassium,  gave  only  84  per  cent 
of  the  total  chlorine. 

A  small  amount  of  potassium  ethylate  usually  separates  out 
during  the  action  of  the  metal,  but  seems  to  have  no  bad  effect 
upon  the  reduction.  After  the  action  has  become  less  vigorous, 
about  2  cm.3  of  alcohol  are  added  and  the  flask  is  carefully  heated 
and  gently  shaken  from  time  to  time  until  the  potassium  is 

*  Ber.  Dtsch.  chem.  Gesell.,  xxxix,  4056. 

t  C.  H.  Maryott,  Am.  Jour.  Sci.,  [4],  xxx,  378. 


448 


METHODS  IN  CHEMICAL  ANALYSIS 


dissolved.  The  contents  are  then  shaken  with  water,  the  water 
layer  is  acidified  with  nitric  acid,  and  the  halogen  precipitated  and 
weighed  as  the  silver  salt.  The  use  of  Volhard's  volumetric 
method  in  estimating  the  halogens  is,  of  course,  feasible,  though 
somewhat  less  accurate.  The  time  required  for  an  analysis, 
exclusive  of  weighings,  is  about  twenty-five  minutes. 
A  series  of  analyses  gave  the  following  results: 

Chlorbenzol,  31.52  Per  Cent  Chlorine. 


Weight  taken, 
grm. 

Per  cent  of 
chlorine  found. 

Per  cent  error. 

0.4036 

31.62 

+0.1 

0-3533 

3I-46 

—0.06 

0.4181 

31.51 

—  o.oi 

0.3278 

31-44 

-0.08 

0.4087 

3I-48 

—0.04 

0.4245 

31-53 

+0.01 

0.3324 

3I-52 

o.o 

Hexachlorbenzol,  74.71  Per  Cent  Chlorine. 


Weight  taken, 
grm. 

Per  cent  of 
chlorine  found. 

Per  cent  error. 

0.1301 
O.I  IQO 

75-39 
75-15 

+0.68 
+0.44 

Brombenzol,  50.1)2  Per  Cent  Bromine. 


Weight  taken, 
gnu. 

Per  cent  of 
bromine  found. 

Per  cent  error. 

0.3976 
0.3921 
0.3928 

5I-29 
51.31 
5I-3I 

+0-37 
+0-39 
+0-39 

p-Chlor aniline,  27.81  Per  Cent  Chlorine. 


Weight  taken, 
grm. 

Per  cent  of 
chlorine  found. 

Per  cent  error. 

0.3081 
0.3456 

27-93 
28.00 

+O.I2 
+0.19 

As  the  hexachlorbenzol ,  brombenzol,  and  p-chloraniline  used 
were  commercial  products  only  once  redistilled  or  recrystallized, 
the  positive  errors  observed  may  be  largely  due  to  impurities. 


CHLORINE;  BROMINE;  IODINE  449 

The   Direct   Determination   of   Chlorine   in    Mixtures   of  Alkali 
Chlorides  and  Iodides. 

Fleischer  justifies  his  use  of  hydrochloric  acid  as  a  standard 
in  alkalimetric  processes  by  his  observation  that  decinormal 
solutions  of  this  acid,  and  even  solutions  of  twice  the  strength 
(7.3  grm.  to  the  liter),  do  not  yield  after  ten  minutes'  boiling 
enough  acid  to  redden  blue  litmus  paper  held  in  the  steam. 
Hydriodic  acid  behaves  similarly  to  hydrochloric  acid  in  the 
matter  of  volatilizing  from  aqueous  solutions ;  but  to  the  decom- 
posing action  of  oxidizing  agents  it  is  far  more  sensitive.  Gooch 
and  Mar  *  have  studied  the  conditions  under  which  hydriodic 
acid  may  be  completely  broken  up  and  iodine  removed  from  the 
solution  by  vaporization  while  the  hydrochloric  is  retained  with- 
out appreciable  loss.  It  was  found  that  the  volatility  of  hydro- 
chloric acid  from  boiling  mixtures  of  potassium  chloride  and 
sulphuric  acid,  dependent  upon  the  concentration  of  the  sul- 
phuric acid  as  well  as  upon  that  of  the  chloride,  is  inconsiderable 
for  i  grm.  of  potassium  chloride  and  5  cm.3  of  concentrated  sul- 
phuric acid,  or  10  cm.3  of  the  [i  :  i]  acid,  in  200  cm.3  of  the  boiling 
solution.  At  300  cm.3  the  dilution  is  sufficient  to  guarantee 
security  against  volatilization  of  hydrochloric  acid  under  the 
conditions  named.  To  remove  the  iodine  of  potassium  iodide 
without  volatilizing  hydrochloric  acid  use  is  made,  for  one 
method,  of  a  ferric  salt  according  to  the  reaction  of  Duflos, 

2  HI  +  Fe2(SO4)3  =  2  FeSO4  +  H2SO4  +  I2, 

and,  for  a  second  method,  the  action  of  a  nitrite  or  of  nitrogen 
trioxide, 

2  KNO2  +  H2SO4  +  2  HI  =  K2SO4  +  2  H2O  +  2  NO  +  I2. 

Use  of  Ferric  In  tests  by  the  first  method  it  was  found  that  from 
Sulphate.  a  volume  of  3OO  cm.3  containing  10  cm.3  of  sulphuric 

acid  [i  :  i],  5  grm.  of  ferric  alum  and  0.005  grm.  of  potassium 
iodide,  every  trace  of  iodine  had  disappeared  so  completely  after 
five  minutes'  boiling  that  nitrous  acid  and  chloroform  collected 
no  color  in  the  cooled  liquid,  but  when  the  amount  of  potas- 
sium iodide  was  increased  to  i  grm.  iodine  was  found  in  con- 
siderable amount  after  boiling  for  an  hour  with  occasional 
replacing  of  water  so  that  the  volume  should  not  decrease  much 

*  F.  A.  Gooch  and  F.  W.  Mar,  Am.  Jour.  Sci.,  [3],  xxxix,  293. 


450 


METHODS   IN   CHEMICAL  ANALYSIS 


below  300  cm.3.  The  reaction  is,  therefore,  plainly  reversible 
under  the  conditions.  When,  however,  a  sufficient  amount  of 
nitric  acid  is  added  to  restore  the  iron  to  the  ferric  state,  boiling 
brings  about  complete  liberation  of  the  iodine.  In  dilute  solu- 
tions the  amount  of  nitric  acid  necessary  to  oxidize  a  fixed  quan- 
tity of  ferrous  salt  is  greater  than  in  concentrated  solutions. 
Thus,  while  o.i  cm.3  of  strong  nitric  acid  should  be  more  than 
enough  to  reoxidize  the  iron  reduced  by  I  grm.  of  potassium 
iodide,  when  the  full  oxidizing  action  is  brought  out,  it  is  neces- 
sary to  add  to  these  dilute  solutions  about  2  cm.3  of  the  acid  to 
complete  the  action  satisfactorily.  The  separation  of  iodine 
and  the  estimation  of  chlorine  according  to  this  process  may  be 
summarized  as  follows:  To  the  solution  of  the  alkali  chloride 
and  iodide  diluted  to  about  400  cm.3,  in  an  Erlenmeyer  flask 
capable  of  containing  a  liter,  are  added  10  cm.3  of  sulphuric  acid 
of  half  strength,  with  2  grm.  of  ferric  sulphate  (either  in  the  form 
of  iron  alum,  or  ferrous  sulphate  oxidized  in  concentrated  solu- 
tion by  about  0.3  cm.3  of  nitric  acid)  and  3  cm.3  of  nitric  acid. 
A  trap  *  is  hung  in  the  neck  of  the  flask,  and  the  liquid  is  boiled 
until  the  steam  which  escapes  no  longer  gives  to  red  litmus  paper, 
after  two  minutes'  exposure,  the  characteristic  gray  blue  due  to 
traces  of  iodine.  Then  I  cm.3  more  of  nitric  acid  is  added  and 
the  test  for  iodine  again  made.  When  no  iodine  is  found  in  the 

Distillation  with  Ferric  Sulphate:   Determination  of  Chloride  in  the  Residue. 


H2S04 
[i  :  i]. 

Fe, 
(S04)3. 

HNO,. 

KC1  =  HC1. 

KI. 

"c3  g 

if 

£-3 

O  w 
<U  -+J 

s.  a 

AgCl  =  HC1 
found. 

Error. 

cm.* 

grm. 

cm.8 

grm.      grm. 

grm. 

.cm.* 

> 
cm.s 

Pi 

grm.       grm. 

grm. 

IO 

2* 

0.4960 

0.2425 

400 

300 

40 

0.9536 

0.2425 

0.0000 

IO 

2* 

0.4970 

0.2429 

400 

300 

40 

0-9534 

o.  2624 

—  O.OOO5 

IO 

2* 

0.4942 

0.2416 

400 

300 

30 

0.9509 

0.2418 

+O.0002 

10 

2* 

2 

0.4969 

0.2429 

.  . 

400 

300 

30 

0-9559 

0.2431 

+O  .  OOO2 

IO 

2* 

3 

0.4956 

0.2423 

400 

350 

30 

0.9546 

o.  2428 

+0.0005 

IO 

2* 

3 

0.4969 

0.2429 

400 

350 

23 

0.9662 

0.2432 

+0.0003 

IO 

2] 

3 

0.4949 

o.  2419 

400 

300 

27 

0.9523 

o.  2422 

+o  .  0003 

10 

at 

5 

0.4970 

0.2429 

400 

250 

55 

0-9559 

0.2431 

+O.OO02 

10 

2t 

5 

0-4955 

o.  2422 

400 

300 

30 

0.9524 

0.2422 

O.OOOO 

IO 

at 

5 

0.4967 

0.2428 

400 

300 

33 

o  .  9546 

o.  2428 

o.oooo 

IO 

at 

6 

0.4964 

0.2427 

400 

300 

30 

0.9550 

0.2429 

+O.OO02 

*  The  iron  was  added  in  the  form  of  iron  alum. 

t  The  iron  was  added  as  FeSO4  oxidized  by  HNOa. 


See  Fig.  6,  page  6. 


CHLORINE;  BROMINE;  IODINE 


451 


escaping  vapor,  silver  nitrate  is  added  in  excess  to  the  contents 
of  the  flask,  the  precipitate  is  settled,  collected  in  a  perforated 
crucible'  on  asbestos,  washed,  dried,  and  weighed  as  silver 
chloride. 

Tests  of  the  method,  with  determinations  in  blank  —  that  is, 
experiments  from  which  the  iodine  was  purposely  omitted  - 
are  detailed  in  the  tabular  statement. 

The  Nitrite  With  pure  sodium  nitrite  at  hand  there  is  probably 

Method.  no  serious  objection  to  introducing  that  substance 

directly  into  the  solution,  but  if  impurities  are  present  it  is  desir- 
able to  generate  the  gas  outside  the  solution.  For  a  generator, 
two  straight  drying  tubes  are  connected  by  a  rubber  tube  and 
set  up  after  the  fashion  of  the  von  Babo  generator,  and  the 
rapidity  of  the  current  is  regulated  to  a  rate  of  five  or  six  bubbles 
to  the  second  by  changing  the  relative  elevation  of  the  gener- 
ator tubes.  The  iodine  separates  immediately  upon  the  intro- 
duction of  the  nitrous  fumes  and  escapes  upon  boiling,  leaving  the 
solution  colorless  in  a  very  short  time.  The  litmus  test  must, 
however,  be  relied  upon  to  indicate  the  removal  of  the  iodine. 
According  to  the  method,  as  developed,  the  solution  of  the 
chloride  and  iodide  contained  in  an  Erlenmeyer  flask  is  diluted 
to  400  cm.3,  10  cm.3  of  sulphuric  acid  of  half  strength  are  added, 
and  the  gas  from  2  grm.  of  sodium  nitrite  acted  upon  by  dilute 
sulphuric  acid  (generated  in  simple  apparatus,  such  as  is  described 
above)  is  passed  with  reasonable  rapidity  into  the  agitated  solu- 
tion. The  liquid  is  boiled  until  colorless,  and  still  further  until 
litmus  paper  placed  in  the  steam  gives  no  reaction  for  iodine 
after  an  exposure  of  two  minutes.  The  contents  of  the  flask  are 


Distillation  with  Nitrite:  Determination  of  Chloride  in  the  Residue. 


NaN02 

H2S04 
[i  :  i]. 

used  in 
gener- 

KC1. =      HC1. 

KI. 

Initial 
volume. 

Final 
volume. 

Time 
in  min- 

AgCl 

found. 

HC1 

found. 

Error. 

ator. 

utes. 

cm.3 

grm. 

grm.        grm. 

grm. 

cm.8 

cm.3 

grm. 

grm. 

grm.' 

10 

2 

0-4953 

0.2421 

I 

400 

350 

2O 

0.9524 

o.  2422 

+  O.OOOI 

10 

2 

0.49750.2432 

I 

400 

350 

16 

0-9573 

0,2434+0.0002 

10 

2 

0.4956 

0.2423 

I 

300 

250 

15 

0.9530 

0.2423      o.oooo 

10 

2 

0-4973 

0.2431 

I 

300 

250 

IS 

0.9550 

o.  2429—0.0002 

10 

2 

0.4964 

0.2427 

I 

300 

250 

15 

0.9550 

0.2429I+0.0002 

IO 

2 

0.4969 

0.2429 

I 

300 

250 

15 

0.9567 

0.2433 

+0.0004 

452  METHODS  IN  CHEMICAL  ANALYSIS 

treated  with  silver  nitrate,  and  the  precipitated  chloride  is 
settled,  collected  on  asbestos  in  the  perforated  crucible,  washed, 
dried  and  weighed. 

The  results  of  test  experiments  are  given  in  the  preceding 
table. 

The  Direct  Determination  of  Bromine  (and  Chlorine)  in  Mixtures 
of  Alkali  Bromides  (and  Chlorides)  with  Iodides. 

The  methods  elaborated  by  Gooch  and  Mar  *  for  the  direct 
determination  of  chlorine  in  mixtures  of  alkali  chlorides  and 
iodides  have  been  studied  by  Gooch  and  Ensign f  with  a  view 
to  similar  application  in  the  determination  of  bromine  in  mix- 
tures of  alkali  bromides  and  iodides.  The  conditions  found 
suitable  in  the  separation  of  chlorine  from  iodine  prove  to  be  in- 
appropriate to  the  separation  of  bromine  from  iodine.  The  use 
of  a  ferric  salt  to  eliminate  iodine  was  not  found  to  be  practi- 
cable under  any  conditions.  Even  at  the  high  dilution  ranging 
from  650  cm.3  to  500  cm.3  bromine  was  likewise  set  free  even 
when  the  concentration  of  the  sulphuric  acid  present  was  very 
low. 

On  the  other  hand  the  nitrite  process,  only  fairly  successful 
when  the  sulphuric  acid  present  is  restricted  to  5  cm.3  of  the  [i  :  i] 
acid  in  a  final  volume  of  500  cm.3,  is  established  as  trustworthy 
when  the  sulphuric  acid  present  is  held  within  the  limits  of  2  cm.3 
to  4  cm.3  of  the  [i  :  i]  mixture.  When  the  quantity  of  sulphuric 
acid  is  still  further  diminished  there  is  evidently  a  slight  tend- 
ency to  show  an  apparent  excess  of  bromide,  due  in  all  probabil- 
ity to  the  retention  of  a  little  combined  iodine  in  the  solution. 
The  best  proportion  for  practical  use  is  probably  3  cm.3  of  the 
half  and  half  acid  in  an  initial  volume  not  less  than  600  cm.3, 
and  this  proportion  proves  to  be  applicable  to  the  separation  of 
iodine  from  an  iodide  associated  with  chloride  as  well  as  bromide. 

The  method  may  be  briefly  summarized  as  follows:  The 
neutral  solution  containing  the  bromide  and  iodide  is  diluted  to 
600  cm.3  or  700  cm.3  (instead  of  400  cm.3,  which  was  found  to  be 
a  sufficient  dilution  in  the  case  of  the  separation  of  chlorine  from 
iodine);  I  cm.3  to  1.5  cm.3  of  strong  sulphuric  acid,  or,  better, 
2  cm.3  to  3  cm.3  of  the  [i :  i]  mixture  (instead  of  the  10  cm.3 

*  See  pages  449,  451. 

t  F.  A.  Gooch  and  J.  R.  Ensign,  Am.  Jour.  Sci.,  [3],  xl,  145. 


CHLORINE;  BROMINE;  IODINE 


453 


employed  in  the  chlorine  separation)  are  added,  a  sufficient 
amount  of  pure  sodium  or  potassium  nitrite  is  introduced  (or, 
if  it  is  preferred,  the  gas  generated  by  the  action  of  dilute  sul- 
phuric acid  upon  the  ordinary  nitrite  and  introduced  from  the 
outside) ;  and  the  liquid  is  boiled,  after  trapping  the  flask,  until 
the  color  has  vanished  and  the  escaping  steam  no  longer  gives 
to  red  litmus  paper  the  color  characteristic  of  iodine.  The 
residual  liquid  is  treated  with  an  excess  of  silver  nitrate  and 
the  precipitated  bromide  filtered  off,  washed,  dried  and  weighed. 
The  process  of  boiling  need  not  extend  beyond  half  an  hour,  or  a 
little  more,  and  care  should  be  taken  that  the  volume  of  the 
liquid  shall  never  be  less  than  500  cm.3.  The  process  has  been 
tested  for  quantities  of  the  potassium  bromide  and  iodide  not 
much  larger  than  0.5  grm.  each.  The  presence  of  0.5  grm.  of 
potassium  chloride  does  not  affect  the  sharpness  of  the  separation. 
The  results  are  given  below : 

Separation  of  Iodine  from  Bromine. 


H2S04 
[i  :  i). 

KI. 

NaNO2 
in  the 
liquid. 

KBr  =  HBr 
taken. 

Initial 
volume. 

Final 
volume. 

Time  in 
min- 
utes. 

AgBr  =  HBr 
found. 

Error  in 
HBr. 

cm.3 

grm. 

grm. 

grm.        grm.. 

cm.* 

cm." 

grin.        grm. 

grm. 

3 

0-5 

0-35 

0.5508 

0.3745 

650 

500 

30 

0.8689 

0-3744 

—  O.OOOI 

3 

o-5 

0-35 

3-5513 

0.3747 

650 

500 

30 

0.8694 

0.3746 

—  O.OOOI 

3 

o-5 

o-35 

0.5513 

0.3747 

650 

500 

30 

0.8699 

0.3748 

+O.OOOI 

3 

0-5 

o-3S 

0.3005 

o  .  2042 

650 

500 

30 

0.4746 

0.2045 

+0.0003 

3 

0-5 

0-35 

0.2759 

0.1875 

650 

500 

30 

0.4358 

0.1878 

+0.0003 

3 

0-5 

i-75 

0.5513 

0.3747 

650 

500 

30 

0.8705 

0.3750 

+0.0003 

3 

o.S 

i-75 

0.55100.3746 

650 

500 

30 

0.8707 

0-3751 

+0.0005 

1 

H2S04 
[I  :i|. 

KI. 

NaN02 
used  in 
gener- 

KBr =  HBr 

taken. 

Initial 
volume. 

Final 

volume. 

Time  in 
min- 

AgBr =  HBr 
found. 

Error  in 
HJr. 

ator. 

utes. 

cm.3 

grm. 

grm. 

grm.        grm. 

cm.3 

cm.3 

grm.        grm. 

grm. 

2 

o-5 

2 

0.5366 

0.3647 

650 

500 

30 

0.84780.3654 

+O.OOO7 

2 

0-5 

2 

0.5369 

0.3650 

650 

500 

30 

0.8472,0.3651 

+  O.OOOI 

2 

o-5 

2 

0.5515 

0.3747 

650 

500 

30 

O.  8687!o.  3742 

—  O.OOO5 

3 

0-5 

2 

0.5371 

0.3652 

650 

500 

30 

0.84590.3644 

—  O.OOOS 

3 

0-5 

2 

0.5365 

0.3647 

650 

500 

30 

0.84650.3647 

O  .  OOOO 

3 

0-5 

2 

0.5368 

0.3649 

650 

500 

30 

0.84860.3656 

+0.0007 

3 

0.5          2 

0.5364 

0.3646 

650 

500 

30 

0.8471  0.3650 

+0.0004 

3 

0-5 

2 

0.5505 

0.3742 

650 

500 

30 

0.86900.3744 

+O.O002 

3 

o.5 

2 

0.0576 

0.0391 

650 

500 

30 

0.09150.0394 

+0.0003 

3 

o-S 

2 

0.0552 

0.0375 

650 

500 

30 

0.08830.0380 

+o  .  0005 

454  METHODS  IN  CHEMICAL  ANALYSIS 

Separation  of  Iodine  from  Bromine  and  Chlorine. 


HfS04 
[i  :  ij. 

KI. 

NaNO2 

KBr. 

HC1. 

Theory 
for 
AgCl+ 
AgBr. 

Found 
AgCl+ 
AgBr. 

Error  in 
silver 
salt. 

Error 
calculated 
asHBr. 

Error 
calculated 
as  HC1. 

cm.3 

grin. 

grin. 

gnu. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm  . 

3 
3 

0.5 

0-5 

0-35 
o-35 

0.5517 
0.55H 

0.4981 
0.4980 

1.8280 
1.8268 

1.8262 
1-8253 

—0.0018 
-0.0015 

—  0.0008 
—0.0005 

—0.0006 
—0.0004 

The  Application  of  lodic  Acid  to  the  Analysis  of  Iodides. 

lodic  acid  may  be  easily  and  completely  reduced  by  an  excess 
of  hydriodic  acid  with  the  liberation  of  iodine  according  to  the 
equation : 

HI03  +  5  HI  =  3  I2  +  3  H20. 

To  apply  this  reaction  to  the  quantitative  estimation  of  iodic 
acid,  it  is  only  necessary  to  add  to  the  free  iodic  acid  or  soluble 
iodate,  in  suitably  concentrated  solution,  an  excess  of  a  solu- 
ble iodide,  to  acidify  —  best  with  dilute  sulphuric  acid  —  and 
to  titrate  with  sodium  thiosulphate  the  iodine  thus  set  free, 
one-sixth  of  the  iodine  found  being  credited  to  the  iodic  acid. 

It  has  been  shown  by  Riegler  *  that  this  reaction  may  be  also 
applied  to  the  quantitative  estimation  of  iodides,  the  iodine  set 
free  upon  the  addition  of  a  known  excess  of  iodic  acid  to  the 
iodide  solution  being  removed  by  petroleum  ether,  and  the 
residual  iodic  acid  determined  as  described  above. 

Gooch  and  Walker  f  have  studied  the  limit  of  applicability  of 
this  reaction  and  have  developed  a  direct  method  for  the  quan- 
titative estimation  of  iodides,  dependent  upon  the  action  of  iodic 
acid  or  an  iodate  in  the  presence  of  free  sulphuric  acid,  neutral- 
ization of  the  solution  by  means  of  an  acid  carbonate,  and  titra- 
tion  of  the  free  iodine  by  arsenious  acid  —  five-sixths  of  the 
iodine  thus  found  being  credited  to  the  iodide  to  be  estimated. 
It  has  been  found  that  by  fulfilling  certain  necessary  conditions 
the  proposed  method  is  entirely  successful,  so  far  as  concerns  the 
estimation  of  iodine  in  iodide  solutions  free  from  large  amounts  of 
chlorides  or  bromides. 

The  degree  of  dilution  at  the  time  when  the  mixture  of  iodide 
and  iodate  is  acidified  has  an  important  influence  upon  the  com- 

*  Zeit.  anal.  Chem.,  xxxv,  305. 

t  F.  A.  Gooch  and  C.  F.  Walker,  Am.  Jour.  Sci.,  [4],  iii,  293. 


CHLORINE;  BROMINE;  IODINE  455 

pleteness  of  the  reaction.  Thus,  the  mean  error  of  test  deter- 
minations in  which  the  volume  at  the  time  of  the  reaction  does 
not  exceed  150  cm.3  is  practically  nothing,  while  the  errors  at 
volumes  of  300  cm.3  and  500  cm.3  amount  to  0.0016  grm.  and 
0.0028  grm.  respectively.  The  doubling  of  the  amount  of  sul- 
phuric acid  used  in  acidifying  does  not  increase  the  amount  of 
iodine  liberated  at  the  highest  dilution.  The  plain  inference 
is  that  the  interaction  between  the  iodide  and  iodate  should  be 
brought  about  in  a  volume  of  liquid  not  much  exceeding  150  cm.3. 
The  apparatus  employed  is  a  reaction  bottle  *  of  500  cm.3  or 
1000  cm.3  capacity,  according  to  requirements,  with  stopcock  and 
thistle-tube  fused  to  the  inlet  tube  and  a  Will  and  Varrentrapp 
absorption  trap  sealed  to  the  outlet.  The  iodide  for  the  test  is 
drawn  from  a  burette  into  the  bottle  and  carefully  washed  down, 
potassium  iodate  in  excess  of  the  amount  theoretically  necessary 
is  added,  and  the  volume  of  the  liquid  is  adjusted  to  150  cm.3. 
The  stopper  with  the  thistle-tube  and  trap  is  put  in  place  and  the 
.trap  is  half  filled  by  means  of  a  pipette  with  a  5  per  cent  solu- 
tion of  potassium  iodide.  Five  cubic  centimeters  of  [i  :  3]  sul- 
phuric acid  are  added  through  the  thistle-tube  and  washed  down, 
the  stopcock  is  closed,  and  the  solution  gently  agitated,  if  neces- 
sary, to  insure  a  complete  separation  of  iodine.  Potassium 
bicarbonate  in  saturated  solution  to  an  amount  about  10  cm.3  in 
excess  of  that  required  to  neutralize  5  cm.3  of  dilute  [1:3] 
sulphuric  acid  is  poured  into  the  thistle-tube  and  allowed  to 
flow  into  the  bottle  slowly  enough  to  avoid  a  too  violent  evolu- 
tion of  gas.  The  stopcock  is  closed  and  the  solution  agitated 
by  giving  to  the  bottle  a  rotary  motion,  at  the  same  time  keep- 
ing the  bottom  pressed  down  upon  the  work-table,  to  prevent  a 
possible  splashing  of  the  iodide  out  of  the  trap  into  the  acid 
solution.  When  the  neutralization  of  the  solution  has  been  com- 
pleted, the  bottle  is  shaken  until  the  last  trace  of  violet  vapor 
has  been  absorbed  in  the  liquid.  The  greater  part  of  the  solu- 
tion in  the  trap  is  then  run  back  into  the  bottle,  the  stopper  is 
removed,  and  the  tube  and  trap  are  carefully  washed,  the  washings 
being  added  to  the  bulk  of  the  solution.  Decinormal  arsenious 
acid  is  introduced  from  a  burette  to  the  bleaching  point,  starch 
emulsion  added,  and  the  solution  titrated  back  with  decinormal 
iodine  (usually  only  a  few  drops)  to  coloration. 
*  Shown  in  Fig.  7,  page  6. 


456 


METHODS  IN  CHEMICAL  ANALYSIS 


The  results  of  experiments  made  in  the  manner  described  upon 
portions  of  a  solution  of  potassium  iodide  standardized  by  the 
arsenate  method  *  are  given  in  the  accompanying  table. 

Analysis  of  Pure  Potassium  Iodide. 


KI  taken, 
grm. 

KI  found, 
grm. 

Error, 
grm. 

0.0814 

0.0816 

+O  .  OOO2 

0.0814 

0.0813 

—  0.0001 

0.0814 

0.0805 

—  O.OOOQ 

0.0815 

o  .  0809 

—  O.OO06 

0.0814 

o  .  0808 

—  O.OOO6 

0.0814 

0.0806 

—  O.OOOS 

0.0814 

0.0812 

—  O.O002 

0.1628 

0.1624 

—  0.0004 

0.1628 

0.1617 

—  o.oon 

0.1628 

0.1621 

—  0.0007 

0.1628 

0.1619 

—  0.0009 

0.1628 

0.1624 

—0.0004 

0.1628 

0.1621 

—0.0007 

0.1628 

0.1626 

—  O.OOO2 

0.2442 

0.2451 

+o  .  0009 

0.2442 

0.2442 

o.oooo 

o  .  2442 

0.2439 

—0.0003 

0.3256 

0.3258 

+O.OOO2 

0.3256 

0.3256 

O.OOOO 

0.3256 

0.3258 

+O  .  0002 

0.3256 

0.3272 

+0.0016 

0.3256 

0.3256 

o.oooo 

0.4071 

0.4076 

+0.0005 

0.4071 

o  .  4080 

+o  .  0009 

0.4071 

0.4073 

+0.0002 

The  presence  of  any  considerable  amount  of  chloride  or 
bromide,  resulting  no  doubt  in  the  formation  of  iodine  chloride 
or  iodine  bromide,  is  prejudicial  to  the  accuracy  of  the  process. 
This  is  shown  in  the  table  following : 

Effects  of  Chloride  and  Bromide. 


KI  taken, 
grm. 

KI  found, 
grm. 

Error, 
grm. 

NaCl  taken, 
grm. 

KBr  taken, 
grm. 

0.0772 

0.0795 

+0.0023 

O.2 

0.0772 

0.0784 

+O.OOT2 

O.2 

0.0771 

0.0823 

+0.0052 

0-5 

0.0773 

0.0819 

+o  .  0046 

o-5 

0.1544 

0.1588 

+o  .  0044 

0-5 

.  .  . 

0.1544 

0.1590 

+o  .  0046 

0-5 

0.0772 

o  .  0802 

+  0.0030 

O.2 

0.0773 

0.0853 

+o  .  0080 

0.2 

0.0772 

0.0873 

+O.OIOI 

.  .  . 

0-5 

0.0772 

0.0861 

+0.0089 

0-5 

0.1544 

o.  1646 

+O.OIO2 

0-5 

0.1543 

0.1626 

+o  .  0083 

0-5 

*  See  page  457- 


CHLORINE;  BROMINE;   IODINE  457 

It  is  plain  that  the  value  of  the  process  in  the  determination 
of  iodine  in  an  iodide  is  restricted  of  necessity  to  those  cases  in 
which  it  is  known  that  chlorides  or  bromides  are  not  present  to 
any  considerable  extent.  For  determining  the  standard  of  a 
solution  of  nearly  pure  potassium  iodide,  employed  in  so  many 
laboratory  processes,  it  is  useful. 

The  lodometric  Determination  of  Iodine  in  Haloid  Salts. 

The  determination  of  iodine  in  a  mixture  of  alkali  chloride, 
bromide  and  iodide  has  been  made  the  subject  of  investigation  by 
Gooch  and  Browning.*  In  this  work  it  was  shown  that  the  iodine 
of  the  iodide  may  be  all  liberated,  under  denned  conditions,  by 
the  combined  action  of  an  arsenate  and  sulphuric  acid,  and  its 
amount  registered  quantitatively  by  the  amount  of  arsenious 
oxide  produced. f  Under  similar  conditions,  the  presence  of  as 
much  as  0.5  grm.  of  sodium  chloride  brings  about  no  formation 
of  arsenious  oxide,  but  does  induce  a  loss  of  that  substance  by 
volatilization  as  arsenic  chloride,  proportionate  to  the  amounts 
of  both  these  substances.  The  effect  of  potassium  bromide  is 
to  produce  trifling  reduction  of  arsenic  .acid  without  volatility. 
Due  correction  of  the  amounts  of  iodine  indicated  by  determi- 
nation of  the  arsenious  oxide  in  the  residue  may  be  made  by  adding 
to  the  indicated  amount  0.008  of  the  product  of  the  weight  of 
chlorine  present  in  chlorides  by  the  weight  of  iodine,  and  sub- 
tracting 0.0024  of  the  weight  of  bromine  in  bromides. 

The  mode  of  proceeding  in  the  determination  of  iodine  in  a 
mixture  of  alkali  chlorides,  bromides  and  iodides,  according  to 
this  method,  may  be  briefly  summarized  as  follows: 

The  substance  (which  should  not  contain  of  chloride  more 
than  an  amount  corresponding  to  0.5  grm.  of  sodium  chloride, 
nor  of  bromide  more  than  corresponds  to  0.5  grm.  of  potassium 
bromide,  nor  of  iodide  much  more  than  the  equivalent  of  0.5  grm. 
of  potassium  iodide)  is  dissolved  in  water  in  an  Erlenmeyer 
beaker  of  300  cm.3  capacity,  and  to  the  solution  are  added  2  grm. 
of  dihydrogen  potassium  arsenate  dissolved  in  water,  20  cm.3  of 
a  mixture  of  sulphuric  acid  and  water  in  equal  volumes,  and 
enough  water  to  increase  the  total  volume  to  100  cm.3,  or  a  little 

*  F.  A.  Gooch  and  P.  E.  Browning,  Am.  Jour.  Sci.,  [3],  xxxix,  188;  xlv,  334. 
t  For  the  reaction,  see  pages  291,  463. 


458 


METHODS  IN  CHEMICAL  ANALYSIS 


more.  A  platinum  spiral  is  introduced,  a  trap  made  of  a  straight 
two-bulb  drying-tube  cut  off  short  is  hung  with  the  larger  end 
downward  in  the  neck  of  the  flask,*  and  the  liquid  is  boiled  until 
the  level  reaches  the  mark  put  upon  the  flask  to  indicate  a  volume 
of  35  cm.3.  Great  care  should  be  taken  not  to  press  the  con- 
centration beyond  this  point  on  account  of  the  double  danger  of 
losing  arsenious  chloride  and  of  setting  up  reduction  of  the  arse- 
nate  by  the  bromide.  On  the  other  hand,  though  35  cm.3  is  the 
ideal  volume  to  be  attained,  failure  to  concentrate  below  40  cm.3 
introduces  no  appreciable  error.  The  liquid  remaining  is  cooled 
and  nearly  neutralized  by  sodium  hydroxide  (ammonia  is  not 
equally  good),  neutralization  is  completed  by  hydrogen  potas- 
sium carbonate,  an  excess  of  20  cm.3  of  the  saturated  solution  of 
the  latter  is  added,  and  the  arsenious  oxide  in  solution  is  titrated 
by  standard  iodine  in  the  presence  of  starch. 


Reduction  of  Ar  senate  and  Determination  of  Arsenious  Oxide  Produced. 


v& 

cm.* 

H2KAsO4. 
grtn. 

NaCl. 

KBr. 

Final 
volume 

cm.8 

Theory  for 
iodine. 

grtn. 

Iodine 
found. 

gnu. 

Error 
found. 

gl'111. 

Error 
corrected. 

2O 

2 

35 

o  .  4080 

0.4079 

—  O.OOOI 

—  O.OOOI 

20 

2 

.  .  . 

35 

.   0.4091 

o  .  4086 

—0.0005 

—0.0005 

20 

2 

35 

0.4083 

o  .  4086 

+o  .  0003 

+0.0003 

2O 

2 

35 

0.0400 

0.0396 

—  O.OOOI 

—0.0004 

2O 

2 

35 

0.0400 

0.0391 

—0.0009 

—0.0009 

20 

2 

35 

0.0400 

0.0400 

o.oooo 

o  .  oooo 

2O 

2 

35 

0:0400 

o  .  0401 

+0.0001 

+O.OOOI 

2O 

2 

35 

0.0040 

0.0037 

—  0.0003 

-0.0003 

2O 

2 

35 

o  .  0040 

0.0038 

—  O.OOO2 

—  O.OOO2 

20 

2 

0-5 

.  .  . 

35 

0.4077 

o  .  4066 

—  O.OOII 

—  O.OOOI 

2O 

2 

0-5 

35 

0.4082 

0.4073 

—0.0009 

+O.OOOI 

2O 

2 

o-5 

35 

o  .  4086 

0.4073 

—  0.0013 

-0.0003 

2O 

2 

0.5 

35 

0.0400 

o  .  0402 

+0  .  0002 

+O.OOOI 

20 

2 

o-5 

35 

0.0400 

0.0395 

—  o  .  0005 

—0.0004 

2O 

2 

o-S 

35 

0.0040 

0.0037 

—  0.0003 

—  O.OOO2 

2O 

2 

o-S 

.... 

35 

0.0040 

0.0037 

—0.0003 

—  O.OOO2 

2O 

2 

0-5 

35 

0.4082 

0.4092 

+0.0010 

+O  .  OOO2 

20 

2 

o.S 

35 

0.4138 

0.4136 

—  O.O002 

—  O.OOIO 

2O 

2 

0-5 

35 

0.4083 

0.4099 

+0.0016 

+0.0008. 

2O 

2 

0-5 

35 

0.0400 

0.0410 

+0.0010 

+O.OOO2 

20 

2 

o-S 

35 

0.0400 

o  .  0404 

-j-o  .  0004 

—0.0004 

2O 

2 

0-5 

35 

o  .  0040 

0.0048 

+0.0008 

0.0000 

2O 

2 

o-S 

35 

0.0040 

o  .  0049 

+0.0009 

+O.OOOI 

2O 

2 

o-5 

o-S 

35 

0.4087 

o  .  4083 

—0.0004 

—  O.OOO2 

20 

2 

o-5 

o-5 

35 

0.4112 

0.4111 

—  O.OOOI 

+0.0001 

2O 

2 

0-5 

o-5 

35 

0.4083 

0.4079 

—  0.0004 

—  O.OOO2 

*  See  Fig.  6,  page  6. 


CHLORINE;  BROMINE;  IODINE 


459 


With  ordinary  care  the  method  is  rapid,  reliable  and  easily 
executed,  and  the  error  is  small.  .  In  analyses  requiring  extreme 
accuracy  all  but  accidental  errors  may  be  eliminated  from  the 
results  by  adding  (algebraically)  to  the  amount  of  iodine  indicated 
an  amount 

i  =  (0.008  X  wt.  Cl  X  wt.  I)  -  (0.0024  X  wt.  Br). 

Results  obtained  by  this  process  are  given  in  uncorrected  and 
corrected  forms. 

Results  *  obtained  in  a  comparison  of  titrations  of  arsenious 
oxide  in  the  residue  with  estimations  of  the  iodine  expelled  under 
the  prescribed  conditions  and  collected  in  the  distillate  show 
close  agreement. 


Comparison  of  A  rsenious  Oxide  in  Residue  and  Iodine  Expelled. 


Iodine 
taken  in 
form  of  KI. 

giro. 

Iodine  found 
from  As2O3 
in  residue. 

gnu. 

Iodine  found 
in  distillate 
by  As2O3. 

grm. 

Iodine  found 
in  distillate, 
by  Na2S203 

grin. 

Error  in 
residue. 

grm. 

Error  in 
distillate. 

grm. 

O  4O^4 

o  40^2 

—  o  0002 

O   4CX7 

O  4O^ 

—  O   OOO2 

o  40^4 

o  40^2 

—  O   OOO2 

o  4054 

o  40^2 

—  O   OOO2 

o  4042 

o  4046 

o  4046 

-|-o  0004 

-|~o  0004. 

o  .  4050 

0  .  405  2 

o  .  4040 

-J-O   OOO2 

—  o  ooio 

o  .  4050 

0.4058 

0.4052 
0.4052 

0.4039 

o  .  405  i 

+0  .  0002 

—  o  0006 

—  O.OOII 

—  o  0007 

o  .  4054 

o  .  4046 

o  .  405  i 

—  0.0008 

—  o  0003 

o  4042 

o  4046 

O   4O3Q 

-j-o  0004 

—  o  ooo  3 

o  .  40  ^  ? 

o  4052 

O   4O<?7 

—  o  0003 

-J-O   OOO2 

The  Determination  of  the  Halogens  by  the  Electrolytic  Reduction  of 
Silver  in  Mixed  Silver  Salts. 

Methods  for  the  estimation  of  silver  in  mixed  silver  salts  of 
the  halogens  have  been  based  by  Gooch  and  Fairbanks  f  on  the 
collection  of  the  precipitated  salts  upon  a  perforated  platinum 
disk  covering  the  asbestos  felt  in  a  perforated  crucible  for  dry- 
ing and  weighing;  the  fusion  of  the  salts  in  contact  with  the 
platinum  disk  to  give  electrical  conductivity;  reduction  of  the 

*  Am.  Jour.  Sci.,  [3],  xlv,  334. 

t  F.  A.  Gooch  and  Charlotte  Fairbanks,  Am.  Jour.  Sci.,  [3],  1,  27. 


460 


METHODS   IN   CHEMICAL  ANALYSIS 


fused  salts  by  making  them  the  cathode  in  the  electrolysis  of  a 
suitable  liquid;  and  the  washing,  igniting  and  weighing  of  the 
reduced  silver. 

saver  chloride  Silver  chloride  and  bromide  are  precipitated; 
and  silver  collected,  washed,  dried  at  150°  C.,  and  weighed  in 
the  filtering  crucible  provided  as  usual  with  a  layer 
of  asbestos  but  in  this  case  covered  with  the  perforated  platinum 
disk.  The  cap  is  put  in  place,  the  crucible  set  upon  an  anvil  to 
keep  it  cool  and  prevent  soaking  of  the  asbestos  with  fused 
silver  salts,  and  the  salts  are  fused  with  a  blowpipe  flame  care- 
fully directed  upon  the  mass  from  above.  A  rubber  band  is 
adjusted  to  cover  the  junction  between  cap  and  crucible.  The 
crucible  is  nearly  filled  with  a  10  per  cent  solution  of  oxalic  acid 
in  25  per  cent  alcohol  and  the  current  passed  in  the  usual  man- 
ner, the  crucible  serving  as  the  negative  electrode.  When  the 
reduction  is  judged  to  be  complete  the  band  and  cap  are  removed, 
the  crucible  set  upon  the  pump,  and  filtration  of  the  liquid  and 
washing  of  the  residue  carried  out  as  usual.  Finally  the  crucible, 
cap  and  residue  are  ignited  at  a  very  low  red  heat  and  weighed. 
The  entire  treatment  is  repeated  until  the  constant  weight  of 
the  residue  shows  that  the  reduction  is  complete.  Results  of 
this  procedure  are  given  in  the  table. 

Electrolytic  Reduction  of  Silver  Chloride  and  Silver  Bromide. 


AgCl  taken, 
gnii. 

AgBr  taken, 
grin* 

Ag  calculated, 
grm. 

Ag  found, 
grm. 

Error, 
grm. 

I.  0608 

o.  7985 

O  .  7990 

+0.0005 

I  4380 

1.0823 

I  .0823 

O.OOOO 

0.9998 

0.7525 

0.7522 

—  0.0003 

0.9959 

0.5721 

0.5723 

-f-O.OOO2 

0.9979 

0.5731 

0.5732 

+O.OOOI 

1.0044 

o  .  4988 

I  .0426 

I  .0422 

—  0.0004 

c-4933 

o  .  4966 

0-6559 

0.6568 

+o  .  oooo 

The  manipulation  of  the  method  is  very  easy,  and  the  results 
show  that  it  is  capable  of  yielding  accurate  results.  In  the 
experiments  recorded  the  current  ranged  from  0.5  to  0.25  amperes, 
and  for  convenience  the  process  was  continued  over  night, 
though  the  reduction  of  amounts  such  as  were  treated  is  usually 
complete  in  six  or  seven  hours. 


CHLORINE;   BROMINE;  IODINE 


461 


Silver  Iodide  by 
Itself  and  in 
Mixture  with 
Silver  Chloride 
or  Silver 
Bromide. 


This  process  which  works  so  well  with  the  mixture 
of  chloride  and  bromide  is  not  applicable  to  the 
reduction  of  silver  iodide  or  to  mixtures  containing 
it.  Experiment  proved  that  the  iodine  set  free  in 
the  electrolysis  works  over  and  over  again  upon  the 
spongy  silver,  constantly  regenerating  silver  iodide  to  a  greater 
or  less  degree.  The  liberated  iodine  may,  however,  be  destroyed, 
without  introducing  anything  objectionable,  by  conducting  the 
electrolysis  in  a  mixture  of  ammonium  acetate,  alcohol,  and 
aldehyde,  —  made  by  neutralizing  two  parts  by  volume  of  ordi- 
nary (40  per  cent)  acetic  acid  with  ammonia,  adding  one  part  of 
ammonia,  one  part  of  alcohol,  and  one  part  of  aldehyde  (75  per 
cent).  Such  a  solution  works  very  well  on  the  whole,  but  as  the 
reduction  progresses  it  frequently  happens  that  a  deposit  of  white 
ammonium  iodate  forms  upon  the  anode  and  introduces  too  great 
resistance  to  the  current.  This  deposit  of  iodate  is,  however,  easily 
removed  from  the  electrode  by  dipping  it  into  hot  water.  When- 
ever the  solution  is  so  exhausted  that  free  iodine  begins  to  appear 
the  liquid  should  be  carefully  decanted  and  replaced  by  fresh  solu- 
tion ;  and  before  the  operation  is  ended  the  decanted  solutions  and 
the  washings  of  the  electrode  should  be  filtered  through  the  cru- 
cible, and  the  residue  submitted  again  to  the  action  of  the  current, 
to  make  it  certain  that  loosened  particles  of  silver  or  silver  salt  pos- 
sibly poured  off  or  removed  on  the  electrode  shall  not  be  lost 


Electrolytic  Reduction  of  Silver  Chloride,  Bromide  and  Iodide. 


AgCl  taken, 
grm. 

AgBr  taken, 
grm. 

Agl  taken, 
grm. 

Ag  calculated, 
grrn. 

Ag  found, 
grrn. 

Error, 
grm. 

0.4779 
o  .  6096 

0.6774 

0.9969 
I  .3703 

0.3596 
0.4588 
0.5098 
0.5727 
o.  7872 

0.3591 
0.4591 
0.5099 
0.5726 
O    78  7  < 

—0.0005 
+0.0003. 
-f-o.oooi 
—  o.oooi 

+o   0003 



.0613 
.0621 
.0140 
.2012 
.  5031 

0.4878 
0.4882 
0.4661 
0.5521 

o  6910 

0.4877 
0.4875 
0.4662 
0.5530 
o  6014. 

—o.oooi 
—  0.0007 

+0.0001 

+0.0009 
+o  0004 

0-5035 

I    OO2O 

0.4984 
o  0008 

0.6653 

I  328^ 

o  .  6653 
I    3283 

o.oooa 

0-4939 

0.5000 

0.6561 
0.5304 

0.6734 
0.5310 

0.6733 
0.5316 

—o.oooi 
+0.0006 

462 


METHODS  IN  CHEMICAL  ANALYSIS 


finally.  The  necessity  of  keeping  the  process  under  occasional  su- 
pervision renders  it  undesirable  to  continue  the  action  over  night. 
The  formation  of  gummy  carbonaceous  matter  not  easily  re- 
moved without  the  application  of  a  degree  of  heat  dangerous  to 
platinum  in  contact  with  silver  was  noted  in  some  cases  of  pro- 
longed action  without  attention.  Many  of  the  experiments  re- 
corded in  the  preceding  table  were  completed  within  seven  hours 
with  a  current  not  exceeding  0.5  ampere. 

These  results  show  that  the  process  affords  an  accurate  reduc- 
tion of  the  chloride,  bromide,  and  iodide  of  silver  and  mixtures  of 
these  salts.  When  the  problem  concerns  the  reduction  of  the 
chloride  and  bromide  only,  preference  is  to  be  given  to  the 
simpler  process  of  reduction  in  alcoholic  oxalic  acid. 


The  Estimation  of  Chlorates  by  Reduction  with  Ferrous  Sulphate. 

Peters  and  Moody  *  have  pointed  out  that  solutions  of  fer- 
rous salts  which  have  been  allowed  to  stand  until  all  dissolved 
oxygen  has  produced  its  effect  are  acted  upon  with  extreme 
slowness  by  atmospheric  oxygen.  Phelps  f  has  made  use  of 
this  mode  of  preparing  ferrous  sulphate  to  test  the  process  sug- 
gested by  Corot,|  as  applied  to  chlorates.  The  results  of  experi- 
ments in  which  the  chlorate  was  treated  with  an  excess  of  approxi- 
mately w/5  ferrous  sulphate  and  15  cm.3  of  sulphuric  acid  [1:3] 
in  a  flask  trapped  to  prevent  mechanical  loss  are  given  below. 

Reduction  by  Ferrous  Sulphate  and  Titration  of  Excess. 


KC1O,  taken. 

Oxygen  value  of 
ferrous  salt  taken. 

Oxygen  value  of 
ferrous  salt  found. 

Error  on  KC1O3. 

grm. 

grm. 

grrn. 

grm. 

0.0500 

0.02756 

0.00814 

—  0.0004 

0.0500 

0.02739 

0.00781 

0.0000 

O.IOOO 

0.04934 

0.01024 

—  O.OO02 

0.1000 

0.04951 

0.01043 

—  O.OOO2 

O.2OOO 

o  .  09086 

0.01247 

+  0.0002 

O.2OOO 

0.09078 

0.01277 

—  0.0008 

0.5000 

0.20552 

0.00993 

—  O.OO06 

O.5OOO 

0.20543 

0.00980 

—  0.0005 

*  Am.  Jour.  Sci.,  [4],  xii,  369;  see  also  page  371. 
t  I.  K.  Phelps,  Am.  Jour.  Sci.,  [4],  xvii,  201. 
J  Compt.  rend.,  cxxii,  449. 


CHLORINE;  BROMINE;  IODINE 

The  mixture  was  brought  to  the  boiling  point,  cooled  to  room 
temperature  by  running  water,  diluted  to  a  volume  of  600  cm.3 
and  titrated  with  potassium  permanganate  after  addition  of 
I  grm.  to  2  grm.  of  manganous  chloride. 

The  lodometric  Estimation  of  Chlorates. 

As  has  been  shown,*  under  conditions  properly  controlled, 
arsenic  acid  in  excess  is  capable  of  expelling  the  iodine  from 
hydriodic  acid  at  the  boiling  temperature  of  the  solution,  being 
itself  reduced  correspondingly  according  to  the  equation 

H3AsO4  +  2  HI  =  H3AsO3  +  H2O  +  I2. 

On  cooling  the  liquid  remaining  after  such  treatment,  and  neu- 
tralizing, the  arsenious  oxide  produced  in  the  reaction  may  be 
reoxidized  iodometrically  in  the  usual  manner,  the  iodine  added 
to  accomplish  this  purpose  being  the  exact  measure  of  the  iodine 
originally  present  as  hydriodic  acid  and  expelled  from  the  acid 
solution  during  the  process  of  boiling. 

Gooch  and  Smith  f  have  found  that  in  a  mixture  of  chloric, 
hydriodic  and  arsenic  acids  the  mutual  action  of  the  chloric  and 
hydriodic  acids  takes  place  according  to  the  equation 

HC1O3  +  6  HI  =  HC1  +  3  H2O  +  3  I2, 

and  goes  steadily  to  completion,  and  that  when  this  effect  is 
accomplished  the  action  of  the  arsenic  acid  in  liberating  iodine 
from  the  residual  hydriodic  acid,  and  in  registering  by  its  own 
reduction  the  amount  of  iodine  thus  set  free,  begins.  The 
reaction  affords,  therefore,  a  method  for  the  estimation  of 
chlorates  which  consists  in  heating  the  chlorate,  in  acid  solution 
and  under  conditions  otherwise  appropriate,  with  a  known 
amount  of  potassium  iodide,  somewhat  in  excess  of  that  theoret- 
ically equivalent  to  the  chlorate,  and  in  presence  of  an  excess  of 
arsenic  acid,  the  arsenious  oxide  produced  in  the  process  being 
determined  iodometrically  and  serving  to  measure  the  amount 
of  iodide  left  undecomposed  by  the  chlorate.  The  difference 
between  the  amount  of  iodide  left  undecomposed  and  that  origi- 
nally introduced  is  the  measure  of  the  chlorate  entering  into  the 
reaction. 

*  See  page  457. 

t  F.  A.  Gooch  and  C.  G.  Smith,  Am.  Jour.  Sci.,  [3],  xlii,  220. 


464  METHODS  IN  CHEMICAL  ANALYSIS 

In  the  practical  application  of  this  process  a  solution  of  approx- 
imately decinormal  potassium  iodide  is  standardized  according 
to  the  method  to  which  reference  has  been  made.  Into  an  Erlen- 
meyer  flask  capable  of  holding  300  cm.3  are  put  a  portion  of  the 
iodide  solution,  2  grm.  of  pure  potassium  arsenate,  20  cm.3  of 
diluted  sulphuric  acid  [i  :  i]  and  enough  water  to  make  the 
entire  volume  a  little  more  than  100  cm.3  A  platinum  spiral  is 
introduced  to  secure  quiet  boiling,  a  trap  made  of  a  straight 
two-bulbed  drying  tube  cut  short  is  hung  with  the  larger  end 
in  the  neck  of  the  flask,*  and  the  liquid  is  boiled  until  the  level 
has  reached  a  mark  upon  the  flask  indicating  a  volume  of  35  cm.3, 
experience  having  shown  that  this  degree  of  concentration  is 
sufficient  and  that  it  is  best  not  to  exceed  it.  The  liquid  remain- 
ing is  cooled  and  nearly  neutralized  by  sodium  hydroxide,  acid 
potassium  carbonate  is  added  to  alkalinity,  20  cm.3  of  a  sat- 
urated solution  of  this  salt  are  added  in  excess,  and  the  arse- 
nious  oxide  in  solution  is  titrated  by  standardized  decinormal 
iodine  in  presence  of  starch.  The  iodine  added  in  the  reoxida- 
tion  of  the  arsenious  oxide  is  taken  as  the  exact  equivalent  of  the 
iodine  expelled  in  boiling. 

Then  a  weighed  amount  of  the  chlorate  is  put  in  an  Erlen- 
meyer  flask,  a  portion  of  the  potassium  iodide  solution,  con- 
taining of  the  iodide  a  weight  amounting  at  least  to  twelve  and  a 
half  times  the  weight  of  the  chlorate  ion  (C1O3),  is  introduced, 
and  the  mixture  is  treated  precisely  as  is  the  iodide  alone  in  the 
process  of  standardization. 

The  difference  between  the  iodine  added  in  this  titration  and 
the  amount  similarly  added  in  the  titration  of  the  standardi- 
zation process  is  the  measure  of  the  chlorate  according  to  the 
reaction 

KC103  +  H2S04  +  6  HI  =  KHS04  +  HC1  +  3  H2O  +  3  I2. 

A  large  excess  of  iodide  over  the  amount  equivalent  to  the 
chlorate  is  unnecessary.  The  amount  of  chloride  produced  from 
the  maximum  weight  of  chlorate  which  may  be  handled  con- 
veniently in  this  process  is  too  small  to  call  for  any  correction  f 
in  the  indicated  amount  of  residual  arsenious  oxide.  Results 

are  shown  in  the  table. 
& 

*  See  Fig.  6,  page  6. 
t  See  page  459. 


CHLORINE;  BROMINE;  IODINE 


465 


Reduction  by  Iodide:   Determination  of  Excess. 


Difference 

vr*\r\ 

H2S04 
[i  :  il 
taken. 

H2KAsO4 
taken. 

KI  =  I 

taken. 

Iodine  cor- 
responding 
to  As2Os 
reduced. 

between 
iodine 
taken  and 
iodine 
added  to 

KC1O8 
taken. 

iiUlUs 
equivalent 
to  differ- 
ence 
between  I 
in  KI  and 

Error. 

oxidize 
As203. 

I  added. 

cm.2 

grm. 

grm.       grm. 

grm. 

grm. 

grm. 

grm. 

grin. 

20 

2 

2.0092 

1.5356 

0.2962 

I  •  2394 

O.2OOO 

O  .  2OOO 

o.oooo 

2O 

2 

2.0092 

1-5356 

0.2973 

1.2383 

O  .  2OOO 

0.1999 

—  O.OOOI 

2O 

2 

1.0380 

0.7934 

0.0570 

0.7364 

0.1185 

O.II88 

+0.0003 

2O 

2 

0.8706 

0.6654 

0.0435 

0.6219 

O.IOOO 

o.  1004 

+0.0004 

20 

2 

0.8706 

0.6654 

0.0429 

0.6225 

O.IOOO 

0.1005 

+0.0005 

20 

2 

0.8706 

0.6654 

0.0435 

0.6219 

O.IOOO 

0.1004 

+o  .  0004 

2O 

2 

0.8706 

0.6654 

0.0435 

0.6219 

O.IOOO 

0.1004 

+o  .  0004 

2O 

2 

0.5023 

0.3839 

0.3208 

0.0631 

O.OIOO 

O.OIO2 

+O.OOO2 

2O 

2 

0.5032 

0.3839 

0.3201 

0.0638 

O.OIOO 

O.OIO3 

+0.0003 

2O 

2 

0.2009 

0.1536 

0.0889 

0.0647 

O.OIOO 

O.OIO5 

+0.0005 

20 

2 

o  .  2009 

0.1536 

0.0903 

0.0633 

O.OIOO 

O.OIO2 

+O.OOO2 

20 

2 

0.2009 

0.1536 

0.0903 

0.0633 

O.OIOO 

O.OIO2 

+O  .  OOO2 

20 

2 

0.1339 

0.1024 

o  .  0405 

0.0619 

O.OIOO 

O.OIOO 

0.0000 

2O 

2 

0.1004 

0.0768 

0.0157 

0.0611 

O.OIOO 

0.0099 

—  O.OOOI 

2O 

2 

o  .  1004 

0.0768 

0.0182 

0.0586 

O.OIOO 

0.0095 

—0.0005 

The  Detection  of  Alkali  Per  chlorates  Associated  with  Chlorides, 
Chlorates  and  Nitrates. 

In  experimenting  at  high  temperatures  with  mixtures  of  alkali 
perchlorates  with  salts  of  the  halogens,  Gooch  and  Kreider  * 
have  succeeded  in  developing  a  simple  and  delicate  method  of 
detecting  perchlorates  associated  with  chlorides,  chlorates  and 
nitrates.  Of  the  various  salts  employed  preference  is  given  to 
fused  zinc  chloride,  chiefly  because,  while  sufficiently  energetic 
in  its  action  upon  the  perchlorate,  it  does  not,  like  manganese 
chloride  or  the  double  chloride  of  aluminium  and  sodium,  evolve 
chlorine  under  the  influence  of  ordinary  air  at  the  high  temper- 
ature of  the  reaction. 

In  making  the  test  the  substance  to  be  examined  is  put  in  dry 
condition  (or  the  solution  of  it  is  evaporated)  in  a  test  tube,  a 
trap  made  by  cutting  off  a  two-bulbed  drying  tube,  or  preferably 
a  mushroom  trap  (Fig.  29),  is  moistened  inside  with  a  solution 
of  potassium  iodide  and  hung  with  the  large  end  downward  in 
the  test  tube,  anhydrous  zinc  chloride  is  added,  and  the  mixture 
*  F.  A.  Gooch  and  D.  Albert  Kreider,  Am.  Jour.  Sci.,  [3],  xlviii,  38. 


466 


METHODS  IN  CHEMICAL  ANALYSIS 


heated  to  fusion.  The  chlorine  evolved  during  the  heating  by 
action  of  the  perchlorate  is  indicated  by  the  iodine  set  free  from 
the  iodide  and  subsequently  washed  with  a  little  water  from  the 
trap  and  tested  with  starch  emulsion.  The  test  for  0.00005  grm- 
of  potassium  perchlorate  is  sure  and  distinct,  as  is  shown  in  the 
experimental  results  given  below: 

Tests  by  Fusion  with  Zinc  Chloride. 


KC1O4  taken, 
grm. 

Indication  by  the 
starch  test. 

O.OOIOO 

Strong. 

0.00050 

Strong. 

O.OOO2O 

Strong. 

O.OOOIO 

Strong. 

O.OOOIO 

Strong. 

0.00005 

Distinct. 

o  .  00005 

Distinct. 

o  .  00003 

Trace. 

0.00003 

None. 

O.OOOOI 

None. 

o.ooooo 

None. 

Obviously,  the  presence  of  an  alkali  chloride  in  the  substance 
examined  cannot  interfere  with  the  certainty  of  action,  but  all 
substances  which  yield  chlorine  by  decomposition  or  by  action 
of  the  air  must  be  removed  before  the  test  is  applied.  To 
break  up  o.i  grm.  of  potassium  chlorate  it  is  only  necessary  to 
treat  with  5  cm.3  of  the  strongest  hydrochloric  acid  and  evapo- 
rate to  dryness.  To  destroy  nitrates  the  plan  of  decomposition 
employed  in  the  quantitative  determination  by  manganous 
chloride  in  hydrochloric  acid  serves  best,*  the  dry  substance  being 
treated  with  2  cm.3  of  the  saturated  solution  of  manganous 
chloride  in  the  strongest  hydrochloric  acid  and  the  liquid  evapo- 
rated. This  method  of  decomposing  the  nitrate  is  peculiarly 
advantageous,  since  the  decomposing  agent  is  itself  an  excellent 
indicator  of  the  completeness  of  the  work  of  removal.  Two  or 
three  treatments  serve  to  remove  the  nitrate  entirely ;  but  before 
proceeding  with  the  test  it  is  necessary  to  remove  the  manganese 
which  has  been  introduced,  inasmuch  as  manganese  chloride  will 
of  itself  evolve  chlorine,  by  exchange  for  oxygen,  when  heated 
in  air.  Sodium  carbonate  in  solution  answers  the  purpose  of 

*  See  page  263. 


CHLORINE;  BROMINE;  IODINE 


467 


removing  the  manganese  (together  with  other  interfering  sub- 
stances) and  the  filtrate  from  the  precipitated  manganous  car- 
bonate leaves  on  evaporation  a  residue,  which,  when  treated  with 
the  anhydrous  zinc  chloride,  gives  indications  for  the  perchlorate 
if  it  is  present  in  appreciable  amount.  A  tenth  of  a  milligram 
of  potassium  perchlorate  may  be  found  with  certainty  when 
associated  with  o.i  grm.  of  the  nitrate,  or  chlorate,  or  both. 
The  results  of  a  series  of  tests  for  potassium  perchlorate  associ- 
ated with  the  chlorate  and  nitrate  of  the  same  element  are 
recorded  in  the  following  table: 

Tests  for  Perchlorate  after  Destruction  of  Chlorate  and  Nitrate. 


KC104  taken, 
grm. 

KClOj  taken, 
grm. 

KNO3  taken, 
grm. 

Indication  of  the 
perchlorate. 

O.OOO5 

Strong.* 

0.0003 

Strong.* 

O.OOO2 

Good.* 

O.OOOI 

Good.* 

O.OOOI 

Trace.* 

o  .  0005 

O. 

O. 

Strong,  f 

0.0003 

O. 

O. 

Good.f 

0.0003 

O. 

0. 

Good.f 

O   OOO2 

O. 

0. 

Trace,  f 

O.OOOI 

0. 

0. 

Trace.f 

o.oooo 

0. 

0. 

None.f 

*  After  procedure  for  the  removal  of  chlorate  and  nitrate  by  HC1. 
t  Chlorate  and  nitrate  removed  by  HC1  +  MnCl2. 


The  lodometric  Determination  of  Perchlorates. 

After  attempting  without  success  to  apply  to  the  quantitative 
determination  of  perchlorates  the  reaction  which  has  been  shown 
above*  to  be  efficient  in  detecting  perchlorates,  Kreider  f  has 
succeeded  in  developing  an  exact  method  which  is  based  upon 
the  iodometric  determination  of  the  oxygen  evolved  from  a 
known  amount  of  perchlorate  on  ignition. 

The  method  is  essentially  the  collection  of  the  oxygen  of  the 
perchlorate;  its  subsequent  passage  into  an  atmosphere  of  nitric 
oxide  over  a  strong  solution  of  hydriodic  acid,  and  the  titration 
of  the  iodine  thus  liberated  with  decinormal  arsenic  in  alkaline 

*  See  page  465. 

t  D.  Albert  Kreider,  Am.  Jour.  Sci.,  [3],  1,  287. 


468  METHODS  IN  CHEMICAL  ANALYSIS 

solution.  A  piece  of  combustion  tubing,  10  or  12  cm.  in  length, 
drawn  out  at  one  end  to  a  narrow  constriction  of  length  sufficient 
to  prevent  the  action  of  the  heat  on  the  rubber  tubing  connect- 
ing it  with  a  receiver,  serves  as  the  heating  chamber.  The  tube 
must  of  course  be  cleansed  from  all  organic  materials  and  cannot 
be  safely  employed  for  more  than  three  fusions.  A  platinum 
boat  (since  porcelain  fuses  to  glass)  serves  for  the  introduction  of 
the  perchlorate  to  the  combustion  tube.  As  a  receiver,  two  lev- 
eling bottles  were  found  vastly  superior  to  a  burette  and  leveling 
tube,  the  glass  stopcocks  of  the  latter  giving  continual  trouble 
by  the  action  of  the  caustic  potash  upon  them.  The  larger 
capacity  of  the  bottle  is  favorable  for  the  reception  of  the  volume 
of  oxygen  evolved  and  its  shape  offers  superior  facilities  for  the 
absorption  of  carbon  dioxide.  Into  the  neck  of  the  receiver  is 
fitted  a  rubber  stopper  carrying  a  capillary  tube  just  even  with 
the  narrower  end  before  the  insertion  of  the  stopper  into  the 
neck  of  the  bottle.  Upon  forcing  in  the  stopper  with  a  slight 
twist,  a  funnel-shaped  orifice  is  made  through  which  the  oxygen 
may  be  withdrawn  without  leaving  any  residue.  The  outer  end 
'of  the  capillary  tube  is  connected  with  the  combustion  tube  by  a 
piece  of  vacuum  tubing  fitted  with  a  pinchcock. 

In  the  process  of  evolving  and  collecting  the  oxygen,  the  per- 
chlorate, weighed  in  the  platinum  boat  and  covered  with  an  equal 
mixture  of  sodium  and  potassium  carbonates,  is  placed  in  the 
combustion  tube  which  is  connected  at  the  larger  end  with  the 
carbon  dioxide  generator.  The  combustion  tube  is  inclined  and 
carbon  dioxide  (obtained  from  a  Kipp  generator  charged  with 
acid  and  marble  previously  boiled  to  expel  air,  and  with  cuprous 
chloride  to  take  up  absorbed  oxygen,  and  washed  first  with  a  solu- 
tion of  iodine  in  potassium  iodide  to  oxidize  traces  of  reducing 
impurities  and  then  with  a  solution  of  potassium  iodide  to  absorb 
volatilized  iodine)  is  sent  through  to  replace  the  air.  Connection 
is  made  with  the  receiver  containing  caustic  potash.  Heat  is 
applied  to  the  combustion  tube  until  the  mixture  in  the  plati- 
num boat  is  in  quiet  fusion  and  the  oxygen  is  drawn  to  the  re- 
ceiver under  slightly  diminished  pressure.  The  tube  is  carefully 
cooled  and  inclined  while  the  current  of  carbon  dioxide  carries  all 
oxygen  to  the  receiver. 

To  bring  about  action  between  the  oxygen  collected  and  hydri- 
odicacid,  through  the  medium  of  nitrogen  dioxide,  a  simple  piece 


CHLORINE;  BROMINE;  IODINE  469 

of  special  apparatus  is  employed.*  A  100  cm.3  pipette  is  cut  off 
short  at  both  ends  and  to  each  end  is  sealed  a  glass  stopcock. 
The  delivery  tube  of  one  of  the  stopcocks  is  cut  off  rather  .short 
after  being  tapered  and  constricted  so  as  to  hold  a  rubber  con- 
nector tightly,  while  the  other  delivery  tube  is  left  long  enough 
to  reach  to  the  bottom  of  an  Erlenmeyer  beaker.  It  is  a  con- 
venience to  have  these  conducting  tubes  3  mm.  or  4  mm.  in  diam- 
eter rather  than  capillaries,  since  for  the  various  connections  all 
air  may  be  expelled  from  them  by  displacement  with  water, 
which  is  easily  accomplished  by  using  a  long-nozzled  wash  bottle. 
By  attaching  the  shorter  end  to  an  ordinary  water  pump  or  to  an 
evacuated  flask,  the  air  is  partially  exhausted.  Then,  the  stop- 
cock is  closed  and  the  bulb  disconnected  and  lowered  into  a 
solution  of  hydriodic  acid  of  approximately  known  strength, 
obtained  by  acidifying  potassium  iodide  (3  grm.  in  30  cm.3) 
with  hydrochloric  acid  immediately  before  use,  to  avoid  liber- 
ation of  iodine  by  the  action  of  air.  When  the  desired  amount 
of  liquid  (30  cm.3)  has  been  drawn  in,  the  stopcock  is  closed  and 
connection  made  with  the  carbon  dioxide,  by  which  all  residual  air 
is  expelled.  Then  the  bulb,  held  so  as  to  prevent  the  escape  of 
the  liquid,  is  again  exhausted  by  attachment  to  the  pump. 
After  admitting  from  a  Kipp  generator,  charged  with  copper 
and  dilute  nitric  acid  kept  hot,  about  10  cm.3  of  nitrogen  di- 
oxide washed  by  passing  through  acidified  potassium  iodide 
and  the  same  reagent  made  alkaline,  attachment  is  made  to 
the  receiver,  and  the  oxygen  is  allowed  to  enter  slowly  under 
diminished  pressure  and  with  continuous  shaking.  The  latter 
precaution  is  essential  to  the  process,  as  otherwise  there  is  im- 
perfect distribution  of  the  hydriodic  acid  and  danger  of  form- 
ing nitric  acid.  But  when  the  solution  of  hydriodic  acid  is 
kept  strong  and  the  shaking  continued  while  the  oxygen  enters 
and  for  a  minute  or  two  afterward,  depending  on  the  rapidity 
with  which  it  was  admitted,  the  oxygen  may  be  allowed  to 
enter  quite  rapidly  without  any  fear  of  imperfect  action. 
The  oxygen  being  immediately  utilized,  the  partial  vacuum 
is  affected  only  by  the  heat  generated,  which  is  scarcely 
noticeable. 

It  is  necessary  of  course  to  prevent  access  of  air  into  the  bulb 
until  the  acid  has  been  neutralized,  to  accomplish  which,  without 

*  See  Fig.  25. 


470 


METHODS  IN  CHEMICAL  ANALYSIS 


loss  of  iodine,  acid  potassium  carbonate  must  be  used,  at  least  for 
the  end  reaction.  To  remove  the  contents  of  the  bulb  for  titra- 
tion,  the  two  delivery  tubes  are  washed  and  filled  with  water, 
the  shorter  end  connected  to  a  supported  funnel  containing  a 
saturated  solution  of  the  acid  carbonate,  and  the  longer  one 
inserted  into  an  Erlenmeyer  beaker  containing  also  a  saturated 
solution  of  the  acid  carbonate  in  amount  sufficient  —  as  previ- 
ously determined  —  to  neutralize  all  the  acid  taken.  On  open- 
ing the  stopcock  of  the  delivery  tube  which  reaches  below  the 
liquid  in  the  beaker  the  bicarbonate  is  drawn  in  by  the  partial 
vacuum,  and  sufficient  carbon  dioxide  is  liberated  to  force  all 
the  liquid  out.  Owing  to  the  consequent  effervescence  as  the 
liquid  gains  its  exit,  the  flow  must  be  regulated  by  the  stopcock 
so  as  to  avoid  loss  of  iodine,  which  is  prevented  by  inclining 
the  beaker  so  that  the  bubbles  strike  quietly  against  its  side. 
To  wash  out  the  bulb,  it  is  raised  almost  horizontally,  so  as  to 
prevent  the  liquid  from  running  through,  and  the  upper  stop- 
cock opened  to  admit  the  bicarbonate  from  the  funnel.  Both 
stopcocks  are  then  closed,  the  bulb  disconnected  and  agitated, 
after  which  it  may  be  washed  with  water  in  presence  of  air 
without  fear  of  liberating  more  iodine.  An  excess  of  deci- 
normal  arsenic  is  then  run  into  the  beaker  and  titrated  back 
with  iodine. 

Results  obtained  by  this  method  are  given  in  the  following 
table. 

lodometric  Determination  of  Evolved  Oxygen. 


BC1O4  taken, 
grm. 

KI  taken, 
grin. 

HC1  taken. 
cm.3 

KC1O4  found, 
grm. 

Error, 
grm. 

O.IOOO 

3 

3 

0.1003 

+0.0003 

O.IOOO 

3 

3 

0.1006 

+0.0006 

O.IOOO 

3 

3 

0.0998 

—  O  .  OOO2 

O.  IOOO 

4 

4 

0.1003 

+o  .  0003 

O.IOOO 

3 

3 

0.1003 

+0.0003 

O.  IOOO 

3 

4 

0.0999 

—  O.OOOI 

O.  IOOO 

3 

3 

0.1003 

+0.0003 

O.  IOOO 

3 

4 

O.IOOI 

+0.0001 

0.1500 

3 

4 

0.1493 

—0.0007 

O.2OOO 

6 

6 

O.I9Q9 

—  O.OOOI 

O.2OOO 

6 

6 

0.2009 

+0.0009 

O.OIOO 

3 

3 

0.0099 

—  O.OOOI 

O.OIOO 

3 

3 

O.OIOO 

o.oooo 

o.oooo 

3 

3 

0.0003 

+0.0003 

CHLORINE;  BROMINE;  IODINE 


471 


The  Estimation  of  Bromates  by  Reduction  with  Ferrous  Sulphate. 

By  treatment  of  a  dissolved  bromate  with  ferrous  sulphate 
standardized  iodometrically,  and  determining  the  residual  ferrous 
salt  similarly,  Phelps  *  has  been  able  to  effect  the  determination 
of  the  oxidizing  power  with  considerable  accuracy.  To  the  solu- 
tion of  the  bromate  is  added  in  a  trapped  flask  an  excess  of  ap- 
proximately n/5  ferrous  sulphate  and  15  cm.3  of  sulphuric  acid 
I1  :  3]-  The  mixture  is  heated  to  boiling,  cooled  in  running  water 
and  nearly  neutralized  with  sodium  carbonate.  Rochelle  salt 
(2  grm.)  is  added  and  an  excess  of  n/io  iodine.  The  mixture  is 
made  alkaline  with  acid  potassium  carbonate  added  in  excess, 
bleached  with  n/io  arsenite  in  presence  of  starch,  and  the 
excess  of  the  arsenite  is  titrated  by  iodine. 

The  following  results  were  obtained: 

Reduction  by  Ferrous  Sulphate  and  Titration  of  Excess. 


KBrO3  taken. 

Oxygen  value  of 
ferrous  salt  taken. 

Oxygen  value  of 
ferrous  salt  found. 

Error  on  KBrO3. 

grm. 

grm. 

grm. 

grm. 

0.0500 

0.01776 

0.00357 

—  O.OOO6 

0.0500 

0.01770 

0.00336 

—  O.OOOI 

o.  1000 

0.03792 

0.00942 

—  0.0008 

O.  IOOO 

0.03792 

0.00922 

—  O.OOOI 

O.  2OOO 

0.06321 

0.00576 

0.0000 

O.  2OOO 

0.06321 

0.00580 

—  0.0002 

0.5000 

0.15670 

0.01342 

—  O.OOI3 

0.5000 

o.  16212 

0.01870 

—  O.OOOS 

The  lodometric  Estimation  of  Bromates. 

Reduction  by  In  connection  with  an  investigation  in  respect  to 
HydriodicAcid.  other  methods  for  the  quantitative  determination 
of  bromates,  Gooch  and  Blake  f  have  studied  the  degree  of  regu- 
larity which  may  be  expected  in  the  method  of  Kratchmer| 
which  is  based  upon  the  action  of  potassium  iodide  in  presence 
of  acid  and  titration  of  the  iodine  set  free  according  to  the  equa- 
tion 

6  HI  +  HBr03  =  HBr  +  3  H2O  +  I2. 

*  I.  K.  Phelps,  Am.  Jour.  Sci.,  [4],  xvii,  201. 

t  F.  A.  Gooch  and  J.  C.  Blake,  Am.  Jour.  Sci.,  [4],  xiv,  285. 

t  Zeit.  anal.  Chem.,  xxiv,  546. 


472 


METHODS  IN  CHEMICAL  ANALYSIS 


The  rate  at  which  this  reaction  proceeds  has  been  investigated 
by  Ostwald,*  by  Noyes  f  and  by  Judson  and  Walker  J.  Time 
of  action,  proportion  of  iodide  to  bromate,  excess  of  acid,  and 
dilution  are  all,  within  limits,  determining  factors  in  the  reac- 
tion; but  in  the  analytical  process  it  is  usually  assumed  that  the 
reaction  goes  soon  to  completion  if  free  acid  and  a  moderate 
excess  of  potassium  iodide  are  present. 

Gooch  and  Blake  show  that  the  reaction  is  very  incomplete 
when  the  excess  of  potassium  iodide  is  only  about  20  per  cent  over 
the  amount  demanded  by  theory,  the  time  of  standing  only  the 
few  minutes  required  for  the  manipulation  of  the  process,  and 
the  dilution  considerable.  When  the  amount  of  potassium 
iodide  used  is  four  times  that  required  by  theory,  the  time  of 
standing  at  least  half  an  hour,  and  the  volume  100  cm.3,  variation 
in  the  amounts  of  acid  above  a  reasonable  minimum  and  in  the 
time  given  to  the  reaction  are  without  apparent  effect. 

Reduction  by  Hydriodic  Acid. 


KBrOj 
taken. 

grm. 

Iodine 
taken. 

grm. 

KI 
taken. 

grm. 

H2S04 
[i  :  i]. 

cm.3 

Time  of 
standing. 

hrs. 

Approxi- 
mate 
volume. 

cm." 

Iodine 
found. 

grm. 

Error  in. 
terms  of 
KBrO,. 

grm. 

o.  1400 

0.6378 

3 

5 

j 

IOO 

0.6343 

—  0.0008 

0.1400 

0.6378 

3 

5 

1 

IOO 

0.6329 

—  o.ooii 

0.1400 

0.6378 

3 

5 

i 

IOO 

0.6329 

—  o.oon 

o  .  1408 

0.6416 

3 

5 

22 

IOO 

o  .  6396 

—0.0004 

0.1408 

0.6416 

3 

5 

22 

IOO 

0.6364 

—  O.OOII 

0.1408 

0.6416 

3 

5 

22 

IOO 

0.6381 

—0.0008 

0.1400 

0.6378 

3 

2-5 

- 

IOO 

o  .  6340 

—0.0008 

0.1400 

0.6378 

3 

2-5 

i 

IOO 

o  .  6336 

—0.0009 

0.1400 

0.6378 

3 

2-5 

1 

IOO 

0.6336 

—0.0009 

0.1400 

0.6378 

3 

i 

1 

IOO 

0-6343 

—0.0008 

0.1400 

0.6378 

3 

o-S 

\ 

IOO 

0.6331 

—  O.OOIO 

HC1 

(sp.  gr.  1.18) 

cm." 

o.  1400 

0.6378 

3 

8 

j 

IOO 

0.6336 

—0.0009 

0.1400 

0.6378 

3 

8 

IOO 

0.6329 

—  O.OOII 

o.  1400 

0.6378 

3 

8 

j 

; 

IOO 

0.6336 

—0.0009 

0.1400 

0.6378 

3 

4 

j 

IOO 

0.6336 

—0.0009 

o.  1400 

0.6378 

3 

4 

-; 

IOO 

0.6333 

—  O.OOIO 

o.  1400 

0.6378 

3 

4 

i 

IOO 

0.6340 

—  0.0008 

0.1400 

0.6378 

3 

4 

I 

IOO 

0.6336 

—0.0009 

o.  1400 

0.6378 

3 

4 

I 

IOO 

0.6333 

—  O.OOIO 

0.1400 

0.6378 

3 

4 

I 

IOO 

0.6336 

—0.0009 

*  Zeit.  phys.  Chem.,  ii.  127.  f  Ibid.,  xix,  599. 

J-  Jour.  Chem.  Soc.,  Ixxiii,  410. 


CHLORINE;   BROMINE;   IODINE  473 

The  results  of  experiments  in  which  measured  amounts  of  stand- 
ard solutions  of  potassium  bromate  (approximately  2.8  grm.  to 
the  liter)  were  thus  treated  with  potassium  iodide  and  hydro- 
chloric or  sulphuric  acid  for  definite  times  in  glass-stoppered  bot- 
tles are  given  in  the  table.  The  iodine  liberated  was  determined 
by  titration  with  sodium  thiosulphate  standardized  against  nearly 
decinormal  iodine  the  value  of  which  was  fixed  by  comparison 
with  decinormal  arsenious  oxide  dissolved  in  acid  potassium 
carbonate. 

The  evolution  of  iodine  by  the  action  of  atmospheric  oxygen 
upon  the  acidified  solution  of  the  iodide  was  found  by  experi- 
ment to  vary  with  the  strength  of  the  acid  and  the  time  of 
exposure  from  o.oooi  grm.  to  0.0003  §rm-  expressed  in  terms 
of  the  bromate,  and  these  values  are  not  greater  than  the  differ- 
ences observed  between  parallel  determinations  of  the  same  sort. 
Probably,  however,  all  the  errors  as  shown  in  the  table  should 
really  be  increased  a  trifle  to  approximate  the  truth,  notably 
those  of  the  experiments  allowed  to  stand  the  longest  period, 
twenty- two  hours. 

The  average  apparent  error  of  the  process  as  applied  to  this 
particular  sample  of  bromate  is  —0.0009  grm.;  anfl  2-5  cm-3  °f 
sulphuric  acid  of  half  strength  [i  :  i]  or  the  equivalent  amount 
of  hydrochloric  acid,  4  cm.3  of  the  acid  of  sp.  gr.  1.18,  in  the 
presence  of  about  3  grm.  of  potassium  iodide,  complete  the  action 
within  half  an  hour,  at  a  dilution  of  100  cm.3,  as  far  as  it  will  go 
under  any  of  the  conditions  tried.  The  phenomenon  noted  by 
Ostwald,*  that  small  amounts  of  hydrochloric  acid  tend  to  force 
the  reaction  more  rapidly  than  equivalent  amounts  of  sulphuric 
acid,  does  not  appear  in  these  experiments,  no  doubt  because  the 
action  was  pushed  to  the  limit  by  the  smallest  amount  of  acid 
employed. 

The  error  of  deficiency  appears  to  be  due  to  impurity  in  the 
sample  of  bromate  rather  than  to  incompleteness  of  the  reaction. 
If  the  reaction  were  incomplete  it  would  be  natural  to  look  for 
the  cause  in  the  possibility  of  the  inhibiting  influence  of  the 
iodine  set  free,  but  it  was  found  in  three  parallel  experiments 
that  the  introduction  of  0.5  grm.  of  free  iodine  dissolved  in  potas- 
sium iodide  failed  to  influence  the  error  appreciably.  In  a 
product  recrystallized  several  times,  and  one  in  which  no  chloride 
*  Loc.  cit.,  page  131. 


474  METHODS  IN  CHEMICAL  ANALYSIS 

can  be  detected,  as  was  the  case  with  the  bromate  of  these  exper- 
iments, the  impurity  most  natural  to  look  for  is  potassium  chlo- 
rate, which  might  resist  removal  in  the  process  of  purification  by 
recrystallization.  The  bromate  was  therefore  tested  by  ignit- 
ing it  and  treating  the  residue  with  potassium  bichromate  and 
sulphuric  acid,  volatilizing  and  collecting  any  chloro-chromic 
anhydride  thus  formed,  and  converting  the  last  into  lead  chro- 
mate.*  Traces  of  chlorine  were  thus  found,  which,  inasmuch  as 
they  were  not  found  in  a  similar  test  upon  the  unignited  bromate, 
must  have  had  their  origin  in  chlorate  intercrystallized  with  the 
bromate. 

Reduction  by  Use  has  been  made  of  the  fact  that  iodic  acid  may 
Arsenious  Acid.  ke  reduced  quantitatively  by  arsenious  acid,f  and 
Gooch  and  Blake  t  have  investigated  the  similar  application  of 
arsenious  acid  to  the  quantitative  estimation  of  bromic  acid 
according  to  the  reaction 

3  H3AsO3  +  HBrO3  =  3  H3AsO4  +  HBr. 

Experiments  made  to  discover  the  limits  within  which  regularity 
of  action  may  be  expected  disclose  the  fact  that  for  conditions 
varying  within  a  rather  wide  range  the  oxidation  of  the  arsenious 
oxide  reaches  a  fairly  definite  limit.  The  bromate  effects  practi- 
cally the  same  proportionate  oxidation  of  arsenious  oxide  in  a 
volume  of  200  cm.3  or  less,  whether  the  sulphuric  acid  of  half 
strength  present  amounts  to  3.5  cm.3  or  10  cm.3,  and  whether 
the  time  of  digestion  is  15  minutes  or  30  minutes  at  the  boiling 
temperature,  one  and  one-half  or  four  hours  on  the  steam  bath,  or 
two  days  at  the  ordinary  temperature.  Setting  aside  experi- 
ments in  which  the  addition  of  acid  did  not  exceed  the  equiva- 
lent of  alkali  carbonate  present  by  more  than  I  cm.3,  the  average 
absolute  variation  from  theory  of  an  entire  list  of  forty- two 
experiments  amounts  to  —0.0007  grin-  in  terms  of  bromate, 
individual  variations  departing  from  the  average  by  about  the 
same  figure.  The  meaning  of  this  fact  seems  to  be  that  the  slight 
error  is  due  to  impurity  in  the  potassium  bromate  employed  and 
not  to  incomplete  oxidizing  action  on  the  part  of  the  bromate.  § 

In  making  use  of  this  process,  the  bromate  is  treated  with  a 
considerable  excess  of  arsenious  oxide  dissolved  in  acid  potassium 

*  See  page  441.  f  See  page  422. 

t  Loc.  cit.,  471.  §  See  lines  4  to  n,  above. 


CHLORINE;  BROMINE;  IODINE 


475 


carbonate,  the  mixture  is  acidified  with  3  cm.3  to  7  cm.3  of  sul- 
phuric acid  [i  :  i],  and  the  liquid,  not  exceeding  200  cm.3  in 
volume,  is  boiled  ten  minutes  or  more,  neutralized  with  acid 
potassium  carbonate  and  titrated  with  iodine.  Results  of  this 
procedure  are  given  in  the  following  table: 

Reduction  by  Ar senile. 


Error  in 

Error  in 

respect  to 

KBrOj 
weighed. 

As203 
taken. 

H2S04 
[i  :  i]. 

Time  in 
minutes. 

As2O3  un- 
changed. 

As?03 
oxidized. 

respect  to 
KBr03 
weighed. 

KBr03  by 
Kratch- 
imer's 

method. 

grm. 

grm. 

cm.s 

grm. 

grm. 

grm. 

grm. 

0.0701 

0.1881 

5 

IO 

O.o66l 

O.I22O 

—0.0014 

—  0.0003 

0.0701 

0.1881 

5 

IO 

0.0650 

0.1231 

—0.0009 

0.0000 

0.0701 

0.2475 

5 

IO 

0.1232 

0.1243 

—  O.OOO2 

+0  .  0007 

0.0701 

0.2475 

5 

10 

0.1236 

0.1239 

—O.OOO4 

+0.0003 

0.0701 

0.2475 

5 

25 

0.1234 

0.1241 

—0.0003 

+o  .  0004 

0.0701 

0.2475 

5 

25 

0.1234 

o.  1241 

—0.0003 

+o  .  0004 

0.1402 

0.4950 

3 

15 

0.2479 

0.2471 

—  O.OOI2 

—  0.0003 

O.  1402 

o.495o 

3 

15 

0.2476 

0.2474 

—  O.OOIO 

—  o.oooi 

o.  1400 

0.6188 

7 

20 

0.3708 

o  .  2480 

—  0.0004 

+o  .  0005 

0.1400 

0.6188 

7 

20 

0.3710 

o  .  2478 

—  0.0005 

+o  .  0004 

0.1400 

0.6188 

7 

20 

0.3706 

0.2482  . 

—  0.0003 

+0.0006 

0.1400 

0.6188 

7 

30 

0.3708 

o  .  2480 

—  O.OOO4 

+o  .  0005 

0.1400 

0.6188 

7 

45 

0.3711 

0.2477 

—  0.0006 

+0.0003 

Here  again,  as  in  the  process  of  reduction  of  the  bromate  by 
hydriodic  acid,  the  results  point  to  a  slight  deficiency  in  the 
oxidizing  power  of  the  bromate.  The  mean  deficiency,  0.0006 
grm.,  differs  from  that  of  the  hydriodic  acid  method  by  about 
0.0003  grm. 

Reduction  by          ^  ^as  Deen  shown  *  that  a  chlorate  may  be  deter- 
Arsenate-iodide  mined  by  adding  to  it  in  solution  potassium  iodide 

Mixture.  •      1  /•  i        « 

in  known  amount,  an  excess  of  an  arsenate,  and  sul- 
phuric acid,  boiling  the  mixture  between  definite  limits  of  con- 
centration, determining  by  titration  with  iodine  the  amount  of 
arsenious  oxide  produced,  and  calculating  the  amount  of  chlo- 
rate present,  from  the  difference  between  the  amount  of  arsenious 
oxide  thus  produced  and  that  which  would  be  produced  if  the 
whole  amount  of  iodide  added  were  allowed  to  act  upon  the 
arsenate  alone.  Gooch  and  Blake  f  have  shown  that  a  bromate 

*  See  page  463. 
f  Loc.  cit.,  p.  471. 


476 


METHODS  IN  CHEMICAL  ANALYSIS 


treated  by  this  process  leaves  a  similar  record  of  its  oxidizing 
power.  According  to  this  process,  the  bromate  is  treated  in  an 
Erlenmeyer  300  cm.3  flask  with  2  grm.  of  potassium  arsenate, 
20  cm.3  of  [i  :  i]  sulphuric  acid  and  water  to  make  the  entire 
volume  a  little  more  than  loocm.3  The  liquid  in  the  trapped 
flask  *  is  boiled  down  to  a  volume  of  35  cm.3,  cooled,  nearly 
neutralized  with  sodium  hydroxide,  and  treated  with  an  excess  of 
20  cm.3  of  acid  potassium  carbonate.  The  arsenious  oxide  in 
solution  is  titrated  by  standardized  n/io  iodine  in  presence  of 
starch.  The  iodine  added  is  taken  as  the  exact  equivalent  of 
iodine  expelled  in  boiling  and  the  difference  between  the  iodine 
thus  expelled  and  that  originally  added  in  the  form  of  potassium 
iodide  is  the  measure  of  the  bromate. 

The  following  table  contains  the  account  of  experiments  made 
in  this  manner  upon  the  sample  of  bromate  the  action  of 
which  in  the  iodide  method  and  in  the  arsenious  acid  method 
is  recorded. 

Reduction  by  Ar senate-Iodide  Mixture. 

H2SO4  [i  :  i]  20  cm.3;   initial  volume  105  to  170  cm.3;   final  volume 
35  cm.3 


KBrO3  taken. 

H2KAsO4 
taken. 

I  value  of  KI 
taken. 

Iodine  corre- 
sponding to 
As203 
produced. 

Iodine  corre- 
sponding to 
KBr03. 

Error  in 
terms  of 
KBr03. 

grm. 

grm. 

grm. 

grm. 

grm. 

grm. 

0.0700 

2 

0.4146 

O  .  0948 

0.3198 

+O.OOO2 

0.0700 

2 

0.4146 

o  .  0954 

0.3192 

0.0000 

0.0700 

2 

0.4146 

0  .  0969 

0.3177 

-0.0003 

0.0700 

2 

0.4146 

0.0975 

0.3171 

—0.0004 

o.  1400 

2 

0.7832 

0.1458 

0.6374 

—  o.oooi 

o.  1400 

2 

0.7832 

0.1463 

0.6369 

—  0.0002 

0.1400 

2 

0.7832 

0.1462 

0.6370 

—  O.OOO2 

The  mean  error  of  these  determinations,  in  which  all  oxi- 
dizing material  is  calculated  as  bromate,  is  not  far  from 
—o.oooi  grm. 

Upon  comparing  these  results  with  those  of  the  iodide  process 
and  of  the  arsenious  acid  process,  it  appears  that  the  deficiencies 
noted  above  are  satisfactorily  accounted  for  by  the  presence  of 
traces  of  chlorate  in  the  bromate.  f 

*  See  Fig.  6,  page  6. 
t  See  page  474. 


CHAPTER  XII. 
MANGANESE;  IRON;  NICKEL;  COBALT. 

MANGANESE. 
The  Determination  of  Manganese  as  the  Sulphate. 

The  estimation  of  manganese  by  the  conversion  of  salts  of 
that  element  with  volatile  acids  to  the  form  of  the  anhydrous 
sulphate  by  the  action  of  an  excess  of  sulphuric  acid,  evapora- 
tion, and  gentle  heating  was  formerly  a  recognized  procedure. 
On  the  authority  of  Rose,*  however,  this  method  was  set  aside 
on  account  of  the  supposed  difficulty  of  removing  the  excess  of 
acid  without  disturbing  the  composition  of  the  normal  salt. 
Oesten,  working  under  Rose's  direction,  found  losses  upon  sub- 
mitting the  crystalline  hydrous  sulphate  MnSO4.5H2O  to  gentle 
ignition  several  milligrams  greater  than  would  have  been  the 
case  if  the  salt  had  been  reduced  to  the  normal  anhydrous  sul- 
phate. At  higher  temperature  the  sulphate  turned  brown  and 
the  losses  were  high.  Volhard  f  studied  the  process  and  showed 
that  manganous  sulphate  may  be  dehydrated,  separated  from  an 
excess  of  sulphuric  acid,  and  brought  into  definite  condition  for 
weighing  as  the  anhydrous  salt  by  careful  and  protracted  heat- 
ing over  a  special  device  of  his  own  —  a  ring  burner  enclosed  in 
a  sheet-iron  casing.  Similar  results  were  obtained  on  evaporating 
with  sulphuric  acid  and  igniting  in  like  manner  an  aqueous  solution 
of  manganous  chloride. 

That  the  process  of  estimating  manganese  in  the  form  of  the 
sulphate  is  both  simple  and  accurate  has  been  shown  by  Gooch 
and  Austin. J  According  to  the  method,  as  described,  the  man- 
ganese salt  of  a  volatile  acid  is  treated  with  sulphuric  acid  in 
amount  more  than  equivalent  to  the  manganese,  the  solution 
is  evaporated  on  the  water  bath  until  the  water  is  removed  as 
far  as  may  be,  and  then,  supported  by  means  of  a  porcelain  ring 
or  triangle,  within  a  larger  porcelain  crucible  used  as  a  radiator, 

*  Ann.  Phys.,  ex,  125. 

t  Ann.  Chem.,  cxcviii,  328. 

J  F.  A.  Gooch  and  Martha  Austin,  Am.  Jour.  Sci.,  [4],  v,  209. 
477 


478 


METHODS  IN  CHEMICAL  ANALYSIS 


so  that  the  bottom  and  walls  of  the  one  are  distant  from  the 
bottom  and  walls  of  the  other  by  an  interval  of  about  I  cm.,  the 
crucible  is  heated  more  strongly.  The  outer  porcelain  crucible 
may  be  heated  over  a  good  Bunsen  flame  to  a  red  heat  without 
risk  of  overheating  the  manganese  sulphate  within  the  inner 
crucible,  and  the  ignition  may  pro'ceed  as  rapidly  as  is  consistent 
with  the  avoidance  of  mechanical  loss  by  spattering. 

Results  of  this  treatment  are  given  in  the  accompanying  state- 
ment: 

Conversion  of  Manganese  Chloride  to  Manganese  Sulphate. 


MnSO4  calculated 
from  AgCl  found 
from  50  cm.3  of 
solution  of  MnCl2. 

MnSO4  found  by 
treatment  of 
50  cm.3  of  solution 
with  H2S04. 

Variation. 

grm. 

grm. 

grin. 

0.3515* 
0.3515* 
0.3515* 

0.3513 
0.3514 
0.3518 

—  O.OOO2 
—  O.OOOI 

+o  .  0003 

*  The  mean  of  two  determinations  —  0.3518  grma.  and  0.3512  grms. 
MnSO4/0tt«d  by  Treatment  of  50  cm.3  of  Each  of  Various  Solutions  with  H2SO4. 


I. 

II. 

III. 

IV. 

V. 

VI. 

<l)  0.3100 
(2)  0.3104 

(3)  0.3096 

(l)  0.3256 
(2)  0.3254 

(l)  0.3534 
(2)  0.3543 

(l)  0.3524 
(2)  0.3520 

(l)  0.3355 
(2)  0.3357 

(l)  0.5475 
(2)  0.5476 

The  Determination  of  Manganese  as  Oxide. 

The  estimation  of  manganese  as  mangano-manganic  oxide, 
Mn3O4,  has  been  so  frequently  criticized  unfavorably  that  the 
method  seems  to  have  passed  from  very  general  use  except- 
ing in  certain  cases  in  which  the  directness  of  the  process  is  a 
temptation  to  incur  the  risk  of  some  uncertainty.  The  produc- 
tion of  the  other  oxides  of  manganese  in  definite  condition  is 
thought  to  be  even  more  uncertain. 

Manganese  dioxide,  MnO2,  begins,  as  Wright  and  Menke  * 

have  shown,  to  lose  oxygen  at  a  temperature  (about  210°  C.)  to 

which  the  hydrated  oxide  must  be  heated  to  free  it  from  water, 

and  very  nearly  that  at  which  the  nitrate  is  converted  into  the 

*  Jour.  Chem.  Soc.,  xxx,  775. 


MANGANESE  479 

dioxide;  so  that  the  chance  of  producing  an  undecomposed 
dioxide  by  the  ignition  of  the  hydra  ted  dioxide  (the  form  in 
which  the  dioxide  generally  appears  in  analytical  processes)  or 
of  the  nitrate  is  small. 

Manganic  oxide,  Mn2O3,  is  produced,  it  is  said,  from  the  other 
oxides  by  ignition  at  a  low  red  heat  under  the  ordinary  condi- 
tions. 

The  manganoso-manganic  oxide,  Mn3O4,  forms,  presumably 
when  an  oxide  of  manganese  is  submitted,  under  ordinary  atmos- 
pheric conditions,  to  the  high  heat  of  the  blast  lamp. 

If  the  proportion  of  oxygen  in  the  surrounding  atmosphere  is 
reduced  below  the  normal,  the  conversion  of  Mn2O3  to  Mn3O4 
goes  on  very  easily,  as  Dittmar  has  shown,*  at  a  temperature 
between  the  melting  points  of  silver  and  aluminum,  while  if  the 
proportion  of  oxygen  in  the  surrounding  atmosphere  is  much 
increased  above  the  normal,  the  reverse  change,  from  Mn3O4  to 
Mn2O3,  tends  to  take  place  at  the  same  temperature.  Gooch 
and  Austinj  have  pointed  out  that  the  condition  most  favor- 
able to  the  production  of  the  oxide  Mn3O4 — a  low  proportion 
of  oxygen  in  the  surrounding  air  —  can  be  maintained  during 
the  ignition  if  the  products  of  combustion  are  made  to  displace 
the  ordinary  air  about  the  crucible,  and  that  this  condition  is 
attained  when  the  crucible  rests  well  within  the  upper  part  of 
the  flame  of  a  large  burner  or  blast  lamp  in  such  manner  that 
the  entire  wall  of  the  crucible  is  covered  by  an  oxidizing  flame. 

If  the  oxide  Mn3O4  in  finely  divided  condition  within  a  plati- 
num crucible  be  moistened  with  nitric  acid  and  heated  gently,  the 
containing  crucible  being  well  above  a  larger  porcelain  crucible 
used  as  a  radiator  and  so  heated  that  only  the  bottom  of  the 
outer  crucible  shows  a  faint  red  glow,  the  oxide  which  remains 
has  approximately  the  composition  of  manganese  dioxide,  MnO2. 

If  the  platinum  crucible  containing  the  oxide  MnO2  be  now 
placed  in  contact  with  the  bottom  of  the  larger  porcelain  cruci- 
ble heated  to  redness  the  oxide  takes  a  composition  approxi- 
mating more  or  less  closely  that  of  manganic  oxide,  Mn2O3. 

If  the  crucible  containing  the  oxide  Mn2Oa  is  now  subjected 
to  the  large  flame  of  a  powerful  Bunsen  burner  or  blast  lamp  in 
such  manner  that  the  oxidizing  flame  covers  the  walls  of  the 

*  Jour.  Chem.  Soc.,  xvii,  294. 

t  F.  A.  Gooch  and  Martha  Austin,  Am.  Jour.  Sci.,  [4],  v,  212. 


480 


METHODS  IN  CHEMICAL  ANALYSIS 


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MANGANESE  481 

crucible,  the  composition  of  the  oxide  remaining  is  generally 
very  closely  that  of  the  oxide  Mn3O4.  The  estimation  of  man- 
ganese in  the  form  of  the  manganoso-manganic  oxide,  Mn3O4,  is 
by  no  means  to  be  considered  utterly  untrustworthy  when  the 
process  is  conducted  in  the  manner  described,  though  it  must  be 
recognized  that  an  irregular  result  may  occur  occasionally.  The 
danger  of  accepting  such  an  irregularity  as  a  correct  indication 
may  be  eliminated  to  a  very  considerable  extent  if  the  precau- 
tion is  taken  invariably  to  moisten  the  ignited  oxide  with  nitric 
acid,  and  ignite  again.  Concordant  results  thus  obtained  may 
be  taken  with  a  fair  degree  of  confidence. 

Results  of  this  procedure,  as  well  as  of  the  treatment  described 
for  producing  in  succession  the  oxides  MnO2,  Mn2O3,  Mn3O4,  are 
given  in  the  preceding  tabular  statement. 

The  Determination  of  Manganese  Separated  as  the  Carbonate. 

Austin  *  has  shown  that  the  precipitation  as  carbonate  in 
presence  of  large  amounts  of  ammonium  chloride,  according  to 
Guyard,f  and  estimation  as  sulphate  |  or  oxide  §  offer  reliable 
means  for  the  determination  of  manganese.  The  manganese 
salt  dissolved  with  a  considerable  amount  of  ammonium  chloride 
(about  10  grm.)  in  200  cm.3  of  boiling  water  is  treated  with 
ammonium  carbonate  in  excess.  The  liquid  is  kept  warm  until 
the  precipitate  subsides  and  is  then  filtered  off  either  upon  paper 
or  upon  asbestos  in  a  perforated  crucible.  The  crucible  and 
precipitate  collected  upon  asbestos  are  ignited  in  the  oxidizing 
flame  of  a  powerful  burner  ||  to  give  the  oxide  Mn3O4  which  is 
weighed  as  such;  the  precipitate  collected  upon  paper  may  be 
ignited  to  the  oxide  and  then  converted  by  treatment  with  a  few 
drops  of  sulphuric  acid  and  gentle  heating  to  the  sulphate. 

Weighing  as  the  carbonate  is  not  feasible;  for,  as  Rose  has 
correctly  stated,  U"  evolution  of  carbon  dioxide  and  oxidation  of 
the  residue  begin  before  the  water  is  thoroughly  removed. 

Experimental  results  are  given  in  the  table. 

*  Martha  Austin,  Am.  Jour.  Sci.,  [4],  v,  382. 

t  Hugo  Tamm,  Chem.  News,  xxvi,  37. 

|  See  page  477. 

§  See  page  478. 

||  See  page  479. 

Tf  Ann.  Phys.,  Ixxxiv,  52. 


482  METHODS  IN  CHEMICAL  ANALYSIS 

Precipitation  as  Carbonate:   Conversion  to  Oxide  and  to  Sulphate. 


NH4C1. 

Mn304 
taken. 

Mn304 
found.     v 

Error. 

MnSO4 
taken. 

MnSO4 
found. 

Error. 

grin. 

grm. 

grm. 

grm. 

grm. 

grm. 

IO 

0.1776 

0.1770 

—0.0006 

IO 

0.1776 

0.1788 

+  O.OOI2 

10 

0.1776 

0.1770 

—  O.OOO6 

10 

0.1776 

0.1774 

—  O.OOO2 



IO 

0.2478 

o  .  2463 

—  O.OOI5 

0.4905 

0.4903 

—  O.OO02 

IO 

O.  II2I 

0.  IIIO 

—  O.OOII 

0.2219 

0.2225 

+O.OOO6 

10 

O.I58l 

0.1584 

+o  .  0003 

0.2128 

0.3126 

—  O.OOO2 

IO 

0.1699 

0.1672 

—  0.0027 

0-3344 

0-3355 

—  0.0009 

The    Determination   of  Manganese    Precipitated    as   Ammonium 

Manganese  Phosphate  and  Weighed  as  Manganese 

Pyrophosphate. 

By  Gibbs'  original  method  manganese  phosphate  was  precipi- 
tated by  hydrogen  disodium  phosphate  in  large  excess  above  the 
quantity  required  to  cause  the  precipitation.  The  flocky  white 
precipitate  was  dissolved  either  in  sulphuric  or  hydrochloric  acid, 
and  precipitated  again  at  the  boiling  temperature  by  ammonia 
in  excess.  This  semi-gelatinous  precipitate,  on  boiling  or  long 
standing  even  in  the  cold,  becomes  crystalline,  the  crystals  form- 
ing beautiful  talcose  scales  which  have  a  pearly  luster  and  a  pale 
rose  color.  The  precipitate  was  filtered  off,  washed  with  hot 
water,  dried  and  ignited.  The  results  obtained  by  Gibbs' 
students  agree  closely  with  the  theory  for  the  pyrophosphate. 

When  a  manganous  salt  is  precipitated  in  the  cold  by  an  excess 
of  an  alkaline  phosphate,  it  falls,  as  Heintz  *  has  shown,  in  the 
form  of  the  trimanganous  phosphate  of  the  formula  Mn3P2O8. 
This  same  phosphate  constitutes  the  greater  part  of  the  precipi- 
tate which  forms  when  a  manganous  salt  reacts  in  the  cold,  in 
the  presence  of  ammonium  chloride,  with  microcosmic  salt  and 
ammonia  in  slight  excess,  but  boiling  or  even  subsequent  stand- 
ing may  effect  a  more  or  less  complete  conversion  of  the  man- 
ganese phosphate  to  the  ammonium  manganese  phosphate. 
The  success  of  the  analytical  process  in  which  manganese  is 
weighed  as  the  pyrophosphate  turns,  therefore,  upon  the  change 
of  the  trimanganous  phosphate,  Mn3P2Os  to  the  ammonium 
*  Ann.  Phys.,  Ixxiv,  449. 


MANGANESE  483 

manganese  phosphate,  NH4MnPO4.  Gooch  and  Austin  *  have 
shown  that  the  presence  of  a  large  amount  of  ammonium  salt 
is  essential  to  the  formation  of  the  precipitate  of  ideal  constitu- 
tion. Apparently  the  proportion  of  ammonium  chloride  present 
to  ammonium  manganese  phosphate  formed  should  be  at  least 
40  :  i,  or,  speaking  approximately,  200  molecules  of  ammonium 
chloride  must  be  present  in  the  liquid  to  every  molecule  of  the 
phosphate  formed ;  and  the  ammonium  chloride  may  be  increased 
almost  to  the  point  of  saturation  of  the  liquid  without  causing 
more  than  a  trifling  solubility  of  the  ammonium  manganese 
phosphate  in  the  presence  of  an  excess  of  the  precipitant.  Fur- 
thermore, the  precipitate  may  be  washed  with  perfect  safety 
with  pure  water  as  well  as  with  slightly  ammoniacal  water,  or 
with  ammoniacal  water  containing  ammonium  nitrate,  if  the 
filtration  is  performed  rapidly  and  the  precipitate  gathered  in 
small  space,  as  is  the  case  when  the  phosphate  is  collected  on 
asbestos  in  the  perforated  crucible.  The  finely  granular  pre- 
cipitate which  may  be  obtained  by  slow  action  of  dilute  ammonia 
added  gradually  to  the  hot  solution  of  the  manganese  salt  appar- 
ently includes  a  portion  of  unconverted  phosphate  which  resists 
the  replacement  of  the  manganese  by  ammonium.  On  the 
other  hand,  the  precipitate  of  flocky  condition  thrown  down  in 
the  cold  passes  easily  to  the  silky  and  crystalline  condition  when 
heated  with  the  proper  amount  of  ammonium  salt,  and  possesses 
a  constitution  approaching  the  ideal  under  such  conditions.  The 
conversion  of  the  flocky  manganous  phosphate  is  so  rapid  that 
the  precipitation  may  be  carried  on  safely  in  glass  vessels.  If 
the  ammonium  chloride  in  the  solution  were  to  be  included  in 
the  precipitate  it  would  volatilize  entirely  during  the  ignition, 
leaving  no  trace  unless,  possibly,  a  portion  of  its  chlorine  were 
to  combine  with  the  manganese.  Tests  for  chlorine  in  the 
residue  of  pyrophosphate  have  resulted  negatively,  no  more 
than  a  mere  trace  being  found  in  any  case,  so  that  the  contam- 
inating effect  of  the  ammonium  chloride  proves  to  be  insignifi- 
cant and  the  responsibility  for  the  excess  of  weight  above 
the  theory  must  apparently  rest  with  the  included  microcosmic 
salt. 

In  the  practical  determination  of  manganese  by  the  phosphate 
method  of  Gibbs,  therefore,  the  presence  of  large  amounts  of 
*  F.  A.  Gooch  and  Martha  Austin,  Am.  Jour.  Sci.,  [4],  vi,  233. 


484 


METHODS  IN  CHEMICAL  ANALYSIS 


ammonium  chloride  is  strongly  advocated.  Good  results  may 
be  obtained  most  easily  and  surely  by  the  following  procedure: 
The  slightly  acid  solution  (in  platinum  or  glass),  containing  in  a 
volume  of  200  cm.3  an  amount  of  manganese  not  more  than  enough 
to  make  0.4  grm.  of  the  pyrophosphate,  20  grm.  of  ammonium  chlo- 
ride, and  5  cm.3  to  10  cm.3  of  a  cold  saturated  solution  of  micro- 
cosmic  salt,  is  precipitated  in  the  cold  by  the  careful  addition  of 
dilute  ammonia  in  only  slight  excess.  The  mixture  is  heated 
until  the  precipitate  becomes  silky  and  crystalline,  the  whole 
is  allowed  to  stand  and  cool  half  an  hour,  the  precipitate  is  col- 
lected upon  asbestos  in  a  perforated  platinum  crucible,  washed 
(best  with  slightly  ammoniacal  water),  dried  at  gentle  heat,  and 
ignited  as  usual.  By  this  process  determinations  of  the  larger 
amounts  of  manganese  —  0.4  grm.  of  the  pyrophosphate  — 
approximate  rather  more  closely  to  the  theoretical  values  than 
do  those  of  the  smaller  amounts  —  0.2  grm.  In  either  case 
the  average  error  should  not  exceed  o.ooio  grm.  in  terms  of 
manganese. 

Results  obtained  by  this  procedure  are  given  in  the  table. 

Determination  as  Manganese  Pyrophosphate. 


Mi^PjO;  equivalent 
to  MnO2. 

Error  in 
terms  of 
Mn2P207. 

grm. 

Error  in 
terms  of 
manganese. 

grm. 

Saturated 
solution  of 
HNaNH4PO<. 

cm.s 

NH4C1. 

grm. 

Total 
volume. 

cm.J 

Manganese 
in  the 
filtrate. 

Taken, 
grin. 

Found, 
grm. 

In  platinum. 


0.1885 

0.1903 

+0.0018 

+0.0007 

5 

20 

200 

None. 

0.1885 

0.1910 

+0.0025 

+O.OOIO 

5 

20 

2OO 

None. 

0.1885 

0.1913 

+0.0028 

+0.001  1 

5 

20 

2OO 

None. 

0.1885 

0.1911 

+0.0026 

+0.0010 

5 

20 

200 

None. 

0.3770 

0.3776 

+o  .  0006 

+O.OOO2 

5 

2O 

2OO 

None. 

0.3770 

0.3773 

+0.0003 

+O.OOOI 

5 

2O 

2OO 

None. 

0.3770 

0.3778 

+0.0008 

+0.0003 

5 

2O 

2OO 

None. 

0.3770 

0-3783 

+0.0013 

+0.0005 

5 

2O 

200 

None. 

In  glass. 


0.1885 

0.1904 

+0.0019 

+0.0007 

5 

20 

2OO 

None. 

0.1885 

0.1898 

+0.0013 

+o  .  0005 

5 

20 

200 

None. 

0.3770 

0.3767 

—0.0003 

—  o.oooi 

5 

2O 

200 

None. 

0.3770 

0.3784 

+0.0014 

+0.0005 

5 

20 

2OO 

None. 

MANGANESE  485 

The  Electrolytic  Determination  of  Manganese. 

For  the  electro-deposition  of  manganese  as  the  dioxide  various 
processes  have  been  described.*  With  stationary  electrodes 
solutions  containing  nitric  acid,  sulphuric  acid,  acetic  acid, 
formic  acid  with  or  without  a  formate,  or  ammonium  acetate 
alone,  with  chrome  alum,  or  with  acetone,  have  been  employed.! 
For  use  with  the  rotating  anode,  solutions  containing  ammonium 
acetate  with  chrome  alum  or  alcohol  have  been  advocated.! 
In  all  these  processes  hydrated  manganese  dioxide  is  deposited 
upon  a  large  anode  which  is  preferably  a  roughened  platinum 
dish  of  considerable  capacity. 

Gooch  and  Beyer  §  have  described  experiments  made  to  test 
the  utility  of  the  electrolytic  filtering  cell  in  the  determination 
of  manganese  as  the  dioxide.  The  procedure  adopted  was  the 
simplest.  Portions,  50  cm.3  each,  of  a  solution  of  pure  man- 
ganous  sulphate,  standardized  by  evaporation  of  measured  por- 
tions and  gentle  ignition  of  the  residue  over  a  radiator, ||  were 
treated,  in  each  case,  with  six  drops  (0.17  cm.3)  of  concentrated 
sulphuric  acid,  and  electrolyzed  in  the  filtering  cell  with  a  current 
of  2  amperes  (N.D.ioo=  5  amp.)  and  20-10  volts,  the  voltage 
decreasing  as  the  solution  became  heated. 

In  one  set  of  experiments  the  process  of  continuous  filtration 
during  electrolysis,  for  which  the  adjustment  of  apparatus  is 
shown  in  Fig.  16,  was  employed. If  The  time  required  for  the 
deposition  of  0.1860  grm.  of  the  dioxide  was  one  hour  and  three- 
quarters.  In  a  second  set  of  experiments  the  closed  cell,  shown 
in  Fig.  15,  was  used  during  the  electrolysis,  and  the  adjustment 
for  filtration  made  subsequently  as  previously  described.**  A 
period  varying  from  two  hours  and  ten  minutes  to  two  hours  and 
fifty  minutes  is  required  in  the  latter  process.  Tests  with  hydro- 
gen dioxide  and  ammonia  showed  that  the  deposition  was  com- 
plete in  the  process  of  continuous  filtration  and  practically  so 
in  the  closed-cell  process.  The  closed-cell  process  naturally 

*  Smith's  Electro-analysis,  edition  of  1907,  page  134  et  seq. 
t  Ibid. 

%  Koester,  Zeit.  Elektrochem.,  x,  553. 

§  F.  A.  Gooch  and  F.  B.  Beyer,  Am.  Jour.  Sci.,  [4],  xxvii,  62. 
||  Am.  Jour.  Sci.,  [4],  v.  209. 
H  See  page  17. 
**  See  page  15. 


486 


METHODS  IN  CHEMICAL  ANALYSIS 


requires  less  attention  during  the  electrolysis,  and  so  it  is  advan- 
tageous to  run  the  process  for  a  period,  perhaps  two  hours,  with 
the  closed  cell,  and  then  to  adjust  the  apparatus  for  filtration 
during  further  electrolytic  action,  in  order  that  floating  particles 
of  the  dioxide  may  be  drawn  to  the  felt  and  completeness  of  pre- 
cipitation may  be  assured.  In  this  way  the  advantage  of  the 
circulating  process  may  be  obtained  with  less  attention  to  manip- 
ulation than  is  required  when  the  filtration  is  continuous  from 
the  start.  The  deposit  is  washed  with  water  after  interruption 
of  the  current,  first  dried  at  200°  for  ten  or  fifteen  minutes  and 
weighed,  and  thereafter  ignited  to  low  redness  in  the  spreading 
flame  of  a  large  burner.*  Results  of  experiments  with  the  cell 
arranged  for  continuous  filtration,  and  of  experiments  in  which 
the  closed  cell  was  used  until  the  electrolysis  was  nearly  over,, 
are  given  in  the  accompanying  table: 

Deposition  of  Manganese  Dioxide. 


Solution 
of  MnS04 
taken. 

cm.» 

2S04 
concen- 
trated. 

cm.8 

Current. 

Theory 
Mn02. 

grm. 

MnO2 
weighed 
as  MnO2. 

grm. 

MnO2 
weighed 
as  Mn3O4. 

grm. 

Error, 
grm. 

Amp.     N.  D.JOO. 

Volt. 

Electrolysis  with  continuous  filtration. 


5° 

0.17 

2 

t: 

2O—  12 

o  1860 

o  1862 

-(-O  OOO2 

0.1858 

—  O  OOO2 

50 

0.17 

2 

C 

2O—  12 

0.1860 

0.1856 

—  o  .  0004 

0.1856 

—  o  .  0004 

5° 

o.  17 

2 

2O~I2 

o  1860 

o  1843 

—  o  0017 

o  1872 

+O  OOI2 

Electrolysis  in  closed  cell  with  subsequent  filtration. 


CO 

0.17 

2 

5' 

2O—  I  2 

o  1860 

o  1860 

o  oooo 

o  I&ZT, 

—  o  0007 

50 

o.  17 

2 

5" 

2O~IO 

o  1860 

o  1856 

—  o  0004 

0.1856 

—  o  0003 

co 

O  17 

2 

1  8—  10 

o  1860 

o  18^3 

—  o  0007 

0.1858 

—  O.OOO2 

The  results  are  evidently  as  good  as  could  be  expected  of  any 
process  which  involves  the  weighing  of  a  manganese  oxide 
brought  to  condition  by  heating.  The  degree  of  oxidation  of 
the  oxide  thrown  down  under  the  conditions  approximates  closely 

*  Am.  Jour.  Sci.,  [4],  v,  214. 


MANGANESE  487 

to  that  of  the  ideal  oxide  represented  by  the  symbol  MnO2.H2O, 
formerly  assigned  by  Riidorff  *  to  the  electrolytically  formed 
oxide,  and  differs  in  that  respect  from  that  of  the  electrolytically 
deposited  oxide  which  was  studied  by  Groeger.f 

The  Determination   of  Manganese  Precipitated   by  the   Chlorate 

Process.. 

Gooch  and  Austin  {  recommend  the  substitution  of  sodium 
chlorate  for  potassium  chlorate  in  the  precipitation  of  manganese 
from  the  nitric  acid  solution,  the  greater  solubility  of  the  sodium 
salt  and  the  consequent  rapidity  with  which  its  decomposition 
takes  place  making  its  use  advantageous. 

According  to  the  treatment  outlined,  the  manganous  nitrate, 
free  from  chlorides  and  sulphates,  is  dissolved  in  concentrated 
nitric  acid  (85  cm.3),  and  treated  with  sodium  chlorate  (5  grm.). 
The  mixture  is  boiled  five  minutes,  more  nitric  acid  (15  cm.8) 
and  a  few  crystals  of  sodium  chlorate  are  added,  and  the  heat- 
ing is  discontinued  as  soon  as  the  liquid  boils  again.  The  in- 
solubility of  the  precipitate,  if  this  procedure  is  followed,  is  so 
great  that  no  more  than  insignificant  traces,  never  exceeding 
o.oooi  grm.,  may  be  recovered  from  the  filtrate  after  filtering  on 
asbestos  upon  the  perforated  cone  or  crucible  and  washing  with 
water.  On  the  other  hand,  prolonged  boiling  after  the  last  addi- 
tion of  chlorate  results  in  considerable  losses  of  manganese  (from 
o.ooio  grm.  to  0.0030  grm.),  due  to  the  reducing  effect  of  lower 
oxides  of  nitrogen  formed  (as  is  always  the  case  in  boiling  nitric 
acid)  after  the  chlorine  dioxide  has  been  expelled.  An  excess  of 
chlorate  at  the  end  of  the  operation  seems  to  be  essential  and  a 
slight  yellow  color  in  the  solution,  due  to  chlorine  dioxide,  is 
a  favorable  indication.  It  is  best  to  filter  the  undiluted  solu- 
tion under  pressure  upon  asbestos  on  a  perforated  cone  having 
a  filtering  surface  of  about  40  cm.2.  Dilution  before  filtration 
tends  to  increase  the  solubility  of  the  manganese. 

While  manganese  is  very  completely  precipitated  in  the  chlo- 
rate process  conducted  with  the  precautions  indicated,  the  condi- 
tion of  oxidation  cannot  be  assumed  to  be  that  of  the  dioxide,  and 
the  indications  of  any  process  which  rests  upon  the  assumption 

*  Zeit.  angew.  Chem.,  1892,  6. 

t  Zeit.  angew.  Chem.,  1895,  253. 

J  F.  A.  Gooch  and  Martha  Austin,  Am.  Jour.  Sci.,  [4],  v,  260. 


488 


METHODS  IN  CHEMICAL  ANALYSIS 


that  the  oxygen  value  of  the  manganese  compound  precipitated 
is  that  of  the  dioxide  is  likely  to  be  erroneous.  If,  therefore,  the 
chlorate  method  is  used  to  separate  manganese,  precautions  must 
be  taken  to  secure  a  definite  condition  of  oxidation  before  apply- 
ing a  process  of  determination  which  depends  upon  the  oxygen 
value  of  the  precipitated  oxide. 

Manganese  Oxide  by  Chlorate  Process. 


Mn  taken. 

Mn  found  upon  the 
hypothesis  that 
MnO,  is  the  oxide 
finally  obtained. 

Error. 

giro. 

grm. 

grm. 

By  reduction  with  potassium  iodide  and  titration 
of  free  iodine  by  thiosulphate. 


o  .  0643 

0.0637 

—  0.0006 

0.0643 

0.0642 

—  o.oooi 

0.0643 

o  .  0642 

—  O.OOOI 

0.0651 

0.0651 

o.oooo 

0.1125 

O.  II2I 

—0.0004 

0.1125 

O.II2I 

—0.0004 

0.1125 

O.  I  I  2O 

—0.0005 

0.1214 

O  .  I  2O6 

—0.0008 

0.1214 

o.  1207 

—0.0007 

o.  1214 

0.1223 

+0.0009 

0.1214 

O.I2I4 

0.0000 

By  reduction  with  arsenious  oxide  and  titration 
of  the  excess  by  iodine  in  presence  of 
Rochelle  salt. 


0.1213 

0.  1212 

—  O.OOOI 

0.1213 

0.  I2OI 

—  O.OOI2 

0.1213 
0.1213 

o  .  i  203 

0  .  I  2O8 

—  o.ooio 

—  O.OOO5 

Gooch  and  Austin  show  that  a  definite  condition  of  oxidation 
may  be  secured  by  dissolving  the  precipitate  in  hydrochloric 
acid,  diluting  a  little,  adding  sulphuric  acid,  evaporating  until 
no  more  hydrochloric  acid  remains,  and  treating  the  manganous 
sulphate  as  follows,  according  to  the  method  of  Wright  and 
Menke:*  The  solution  of  manganous  sulphate  (not  exceeding 
0.5  grm.),  very  nearly  neutralized  by  potassium  carbonate,  is 
*  Jour.  Chem.  Soc.,  xxxvii,  36. 


NICKEL 


489 


mixed  with  a  solution  of  zinc  sulphate  (2  grm.),  and  a  freshly 
made  and  carefully  filtered  dilute  solution  of  potassium  perman- 
ganate (1.5  grm.)-  The  liquid  now  amounting  to  about  500  cm.3 
is  heated  to  80°  and  acid  potassium  carbonate  is  added  in  quan- 
tity a  little  more  than  enough  to  neutralize  the  free  acid.  The 
loose  precipitate  is  collected  upon  asbestos  and  after  careful 
washing  is  returned  to  the  flask  in  which  the  precipitation  was 
made.  The  oxygen  value  of  the  oxide  thus  obtained,  now 
that  of  the  oxide  MnO2,  may  be  determined  by  any  appropriate 
method  of  treatment. 

The  preceding  table  contains  the  results  of  two  different 
methods  for  determining  the  oxygen  value  of  manganese  dioxide 
separated  by  the  chlorate  method  and  brought  to  ideal  condition 
by  the  procedure  described. 


NICKEL  (COBALT). 
The  Electrolytic  Determination  of  Nickel  with  the  Rotating  Cathode. 

Gooch  and  Medway*  determine  nickel  by  deposition  upon 
the  rotating  cathode, f  from  50  cm.3  of  solution  containing  one- 
half  its  volume  of  concentrated  ammonia  and  a  gram  of  am- 
monium sulphate.  It  should  be  especially  noted  that  the  solution 
must  be  kept  within  the  limit  of  volume  indicated  above,  as 
further  dilution  lengthens  the  time  necessary  for  complete 
deposition. 

Deposition  of  Nickel. 


Nickel  taken.^ 
gnu. 

Nickel  found, 
grm. 

Error, 
grm. 

Current, 
amp. 

N.D.100. 

Time, 
min. 

0.0954 

0.0954 

0.0000 

i-5 

5 

30 

0.0954 

0.0953 

—  o.oooi 

3 

10 

25 

0.0954 

0.0956 

+0.0002 

3 

10 

25 

0.0954 

0.0953 

—o.oooi 

3-5 

11.7 

20 

0.0954 

0.0955 

+O.OOOI 

3-5 

ii.  7 

2O 

0.1738 

0.1736 

—  O.OOO2 

3-5 

11.7 

25 

0.1738 

0.1740 

+0.0002 

3-5 

11.7 

25 

0.1738 

o  .  i  740 

+0.0002 

4 

13-3 

25 

0.1738 

0.1737 

—  O.OOOI 

4 

13-3 

25 

0.1738 

0.1738 

o.oooo 

4 

13-3 

25 

*  F.  A.  Gooch  and  H.  E.  Medway,  Am.  Jour.  Sci.,  [4],  xv,  323. 
t  See  page  n. 


490 


METHODS  IN  CHEMICAL  ANALYSIS 


The  Estimation  of  Nickel  by  Precipitation  as  the  Oxalate  and 
Titration  with  Potassium  Permanganate. 

Classen  *  has  shown  that  nickel  may  be  completely  precipi- 
tated by  treating  the  soluble  nickel  potassium  oxalate  with  a 
large  amount  of  acetic  acid,  and  estimated  by  igniting  the  oxalate 
and  weighing  the  oxide.  Ward  f  has  adapted  this  process  to  the 
volumetric  determination  of  nickel  by  titration  of  the  precipi- 
tated oxalate  with  permanganate.  Finding  that  the  precipitate 
formed  by  precipitation  with  potassium  oxalate  tends  to  include 
the  alkali  oxalate  and  that  the  nickel  oxalate  thrown  down  from 
the  boiling  acetic  acid  solution  of  a  nickel  salt  falls  in  finely 
divided  condition  and  is  difficult  to  filter,  Ward  makes  the  first 
precipitation  in  water  solution  and  adds  the  acetic  acid  after- 
ward. The  nickel  salt  is  dissolved  in  a  definite  amount  of  water, 
crystallized  oxalic  acid  is  added  to  the  boiling  solution,  and  an 
equal  volume  of  acetic  acid  is  added.  The  precipitate  is  allowed 
to  settle  over  night,  filtered  on  asbestos  in  a  perforated  crucible, 
and  washed  with  small  amount  of  water.  The  crucible  is  placed 
in  a  beaker  containing  about  25  cm.3  of  dilute  [1:3]  sulphuric 
acid,  and  heat  is  applied  to  dissolve  the  oxalate.  The  vol- 
ume of  the  solution  is  made  up  to  about  200  cm.3  with  water, 
and  cobalt  sulphate  is  added  until  the  green  color  of  the  nickel 
salt  is  neutralized  and  a  slight  pink  tinge  appears.  The  addition 
of  the  cobalt  salt,  recommended  by  Gibbs,  is  necessary  to  secure 
a  definite  end  point.  The  solution  is  heated  to  boiling  and  the 
titration  is  made  in  the  usual  manner  with  permanganate.  Re- 
sults of  this  procedure  are  given  in  the  table. 

Precipitation  by  Oxalic  Acid. 


Nickel  taken 
as  the 
sulphate. 

Volume  of 
water  solution 
at  precipitation. 

Oxalic 
acid. 

Acetic  acid 
added. 

Nickel  found. 

Error. 

grm. 

cm.* 

grm. 

grin. 

grm. 

grm. 

0.0050 

IOO 

2 

IOO 

0.0054 

+0.0004 

0.0257 

100 

2 

IOO 

0.0258 

+O.OOO7 

0.0503 

IOO 

2 

IOO 

O.O5O2 

—  O.OOOI 

0.1257 

IOO 

2 

IOO 

O.I27I 

+0.0014 

*  Zeit.  anal.  Chem.,  xvi,  470. 

t  H.  L.  Ward,  Am.  Jour.  Sci.,  [4],  xxxiii,  336, 


NICKEL 


491 


The  Detection  of  Nickel  in  Presence  of  Cobalt. 

Browning  and  Hartwell  *  have  modified  advantageously  the 
method  of  Clarke  f  for  the  separation  of  nickel  and  cobalt, 
according  to  which  cobalt  precipitated  as  the  ferricyanide  re- 
mains insoluble  in  ammonium  hydroxide  while  the  nickel  salt 
is  dissolved. 

The  method  as  modified  may  be  described  as  follows:  To  not 
more  than  o.i  grm.  of  the  salts  of  the  two  elements  in  about 
5  cm.3  of  water  a  few  drops  of  a  saturated  solution  of  alum  are 
added,  free  mineral  acid  is  neutralized  with  ammonium  hydroxide, 
and  the  solution  is  made  faintly  acid  with  acetic  acid.  To  this 
solution  is  added  about  0.5  grm.  of  solid  potassium  ferricyanide 
with  shaking  to  effect  the  solution  of  this  reagent  and  the  pre- 
cipitation of  the  nickel  and  cobalt  salts.  The  precipitate  is 
treated  with  about  5  cm.3  of  strong  ammonium  hydroxide  and 
the  mixture  is  filtered.  To  the  filtrate,  which  should  have  no 
reddish  color,  a  piece  of  sodium  or  potassium  hydroxide  about 
the  size  of  a  pea  is  added  and  the  mixture  is  boiled.  The  appear- 

Detection  of  Nickel. 


CoS04.7H20 

gnu* 

NiS04.7H20 

grm. 

KA1(S04)2 
saturated 
solution. 

cm.» 

KjFeC.N,. 
grm. 

NH4OH 
(concen- 
trated). • 

cm.» 

NaOH, 
solid. 

grm. 

Result. 

.... 

O.OIOO 

2 

0-5 

5 

<    I 

Heavy 
precipitate. 

— 

O.OO5O 

2 

0-5 

5 

<    I 

Heavy 
precipitate* 

.... 

0.0010 

0.0003 

2 
2 

o-S 
o-S 

5 
5 

<    I 

Heavy- 
precipitate. 
Distinct. 

.... 

O.OOOI 

2 

o.S 

5 

Plain. 

O.  IO 

2 

o.  <c 

c 

None. 

0.10 

O.OIOO 

2 

O 

0-5 

O 

5 

Heavy. 

0.10 

0.0050 

2 

0.5 

5 

Distinct. 

0.10 

0.0030 

2 

o-5 

5 

Very  faint. 

0.05 

2 

o-5 

5 

None. 

0.05 

O.OIOO 

2 

o-S 

5 

Heavy. 

0.05 

0.0050 

2 

o-5 

5 

Distinct. 

0.05 

0.0030 

2 

0-5 

5 

Plain. 

0.05 

0.0010* 

2 

0.5 

5 

Faint. 

*  Equivalent  to  0.0002  of  the  metal. 


*  Philip  E.  Browning  and  John  B.  Hartwell,  Am.  Jour,  Sci.,  [4],  x,  316. 
t  F.  W.  Clarke,  Am.  Jour.  Sci.,  [3],  xlviii,  67. 


492 


METHODS  IN  CHEMICAL  ANALYSIS 


ance  of  black  nickelic  hydroxide,  very  small  amounts  showing 
first  as  a  dark  coloration,  indicates  nickel. 

The  preceding  table  records  the  experimental  results. 

The  Separation  of  Nickel  and  Cobalt  by  the  Etherial  Solution  of 
Hydrochloric  Acid. 

The  work  of  Havens  *  shows  that  very  small  amounts  of  nickel 
may  be  separated  from  cobalt,  both  taken  as  the  chlorides,  by 
treating  the  dry  salts  with  the  least  amount  of  water  which  will 
effect  solution  (i  cm.3),  adding  ether  (10  cm.3  to  15  cm.3)  and 
saturating  the  mixture,  cooled  in  running  water,  with  gaseous 
hydrochloric  acid.  When  saturation  is  complete  the  precipitated 
nickel  chloride  is  filtered  upon  asbestos  in  the  perforated  crucible 
and  washed  with  ether  previously  saturated  with  hydrochloric 
acid.  From  the  etherial  solution  containing  the  minimum  of 
water  the  precipitation  of  nickel  chloride  is  practically  complete, 
but  when  the  amount  of  cobalt  present  exceeds  a  few  milli- 
grams errors  due  to  inclusion  of  cobalt  chloride  in  the  mass  of 
precipitated  nickel  chloride  appear. 

Results  of  this  method  of  separation  are  given  below.  The 
nickel  in  the  precipitated  chloride  and  the  cobalt  in  the  filtrate 
were  determined  by  the  electrolytic  process  in  the  experiments 

recorded. 

Separation  by  Ether-Hydrochloric  Acid. 


Nickel  taken 
as  the  hydrous 
chloride. 

Nickel  found. 

Error. 

Cobalt  taken 
as  the  hydrous 
chloride. 

Cobalt  found. 

Error. 

0.0068 

o  0066 

—  O  OOO2 

O.OOQO 

0.0090 

O   OOOO 

o  0300 

O.OOQO 

o  .  0469 

0.0091 
o  .  0490 

+O.OOOI 
-J-O.OO2I 

0.0123 

0.0700 

0.0127 

+O.OOO4 

o  .  0468 

o  o<;o3 

-j-o  oo3<? 

o  0700 

0.0472 

0.0493 

+O.OO2I 

0.0700 

IRON. 

The  Determination  of  Iron  in  the  Ferric  State  by  Reduction  with 

Sodium  Thiosulphate  and  Titration  of  the  Excess  of  the 

Latter  with  Iodine. 

The  action  of  sodium  thiosulphate  upon  a  ferric  salt  takes 
place  with  the  formation  of  a  ferrous  salt  and  a  tetrathionate, 
according  to  the  following  expression 

2  FeCl3  +  2  Na2S2O3  =  2  FeCl2  +  Na2S4O6  +  2NaCl. 
*  Franke  Stuart  Havens,  Am.  Jour.  Sci.,  [4],  vi,  396. 


IRON  493 

After  many  attempts  by  others  *  to  utilize  the  reaction  in  the 
quantitative  estimation  of  iron,  Norton  f  has  succeeded  in  plac- 
ing the  method  upon  a  successful  basis  for  the  determination  of 
moderate  amounts  of  ferric  iron  by  direct  titration.  Norton  has 
studied  carefully  the  possible  sources  of  error:  incompleteness  in 
the  reduction  of  the  ferric  salt ;  decomposition  of  the  thiosulphate 
by  the  acid,  resulting  in  the  subsequent  over-run  of  iodine;  the 
possible  tendency  of  the  ferric  salt  under  concentration  to  oxidize 
the  thiosulphate  to  the  condition  of  the  sulphate  rather  than  to 
that  of  the  tetrathionate ;  and  finally  the  oxidizing  action  of 
the  air,  which  may  tend  to  keep  up  progressive  oxidation  of  the 
iron  salt  and  excessive  expenditures  of  thiosulphate.  The  first 
three  sources  of  difficulty  tend  to  produce  errors  of  deficiency; 
the  fourth  an  error  of  excess. 

It  is  shown  that  with  quantities  of  ferric  oxide  present  up  to 
o.i  grm.  the  dilution  may  vary  from  400  cm.3  to  1000  cm.3  for 
each  cubic  centimeter  of  strong  hydrochloric  acid  and  still  give 
excellent  results.  At  a  dilution  greater  than  1000  cm.3  the  action 
of  the  thiosulphate  is  incomplete  in  presence  of  I  cm.3  of  the  acid, 
and  at  a  smaller  dilution  than  400  cm.3  the  decomposing  action 
of  the  acid  on  the  thiosulphate  becomes  noticeable.  When  larger 
quantities  of  iron  .oxide  are  dealt  with,  the  dilution  must  be  in- 
creased proportionately  to  the  quantity  of  ferric  oxide.  On  this 
account  it  is  necessary,  assuming  that  the  quantity  of  acid  present 
is  always  kept  within  the  maximum  strength  mentioned,  I  cm.3 
to  400  cm.3,  to  regulate  the  dilution  by  the  approximate  quan- 
tity of  the  iron  so  that  not  less  than  400  cm.3  of  water  shall  be 
used  to  every  o.i  grm.  of  iron  oxide  present.  Under  properly 
regulated  conditions  of  dilution  as  regards  acid  and  the  iron  salt, 
the  reduction  is  completed  in  from  five  to  ten  minutes. 

Under  the  conditions  of  acidity  and  dilution  laid  down,  the 
process  of  reduction  is  complete  at  the  ordinary  room  tempera- 
ture within  ten  minutes  after  the  introduction  of  the  thiosulphate, 
and  experience  shows  clearly  the  danger  of  submitting  mixtures 
of  sodium  thiosulphate  and  acid  to  temperatures  much  above  the 

*  Sherer;  Gelehrte  Anzeigen  der  konig.  Bayrisch.  Acad.,  1859.  Mohr, 
Ann.  Chem.,  cxiii,  269.  Kremer  and  Landolt,  Zeit.  anal.  Chem.,  i,  214. 
Oudemanns,  Zeit.  anal.  Chem.,  vi,  129;  ix,  362.  Haswell,  Repert.  anaL 
Chem.,  i,  179.  Bruel,  Compt.  rend.,  xcvii,  954. 

f  John  T.  Norton,  Jr.,  Am.  Jour.  Sci.,  [4],  viii,  25. 


494 


METHODS  IN  CHEMICAL  ANALYSIS 


ordinary.  On  the  other  hand,  artificial  reduction  of  temperature 
tends  to  retard  the  action  to  an  undesirable  degree. 

To  complete  the  reduction  within  a  reasonable  time  under  the 
conditions  of  acidity  and  concentration  there  should  always  be 
present  an  excess  of  thiosulphate  in  amount  from  15  cm.3  to 
35  cm.3  of  the  n/io  solution. 

The  following  procedure  is  recommended:  The  dilution  must 
be  at  least  400  cm.3  for  each  o.i  of  a  grm.  of  iron  oxide  pres- 
ent; the  quantity  of  acid  should  never  exceed  I  cm.3  of  the 
strong  acid  to  each  400  cm.3  of  water;  the  time  of  reduction  must 
be  short  to  avoid  progressive  oxidation;  the  temperature  of  the 
solution  should  be  kept  at  the  normal  temperature  of  the  atmos- 
phere; and  the  excess  of  sodium  thiosulphate  present  should 
never  be  less  than  15  cm.3  of  the  n/io  solution.  In  the  case  of 
large  dilution  freshly  boiled  water  should  be  used. 

Titration  by  Thiosulphate. 


Fe203 
taken. 

grm. 

Fe20j 
corrected. 

grm. 

Dilution. 
cm.» 

HC1 
cone. 

cm.8 

Excess 
Na2S20,. 

cm.1 

Fe203 
found. 

grm. 

Error, 
grm. 

0.0125 

0.0125 

200 

1 

23-5 

0.0125 

O.OOOO 

0.0250 

0.0250 

400 

i 

21.98 

0.0250 

O.OOOO 

0.0250 

0.0250 

400 

% 

17 

O*.0250 

O.OOOO 

0.0250 

0.0250 

400 

^ 

17 

0.0250 

0  .  0000 

0.0500 

o  .  0499 

400 

i 

24 

o  .  0498 

—  O.OOOI 

o  .  0500 

o  .  0499 

400 

i 

19 

o  .  0498 

—  O.OOOI 

O.O5OO 

0.0499 

400 

i 

I5-I 

0.0497 

—  O.O002 

o  .  0500 

o  .  0499 

400 

1 

19 

o  .  0498 

—  O.OOOI 

O.  IOOI 

0.0999 

400 

23.1 

0.0993 

—  0.0006 

O.IOOI 

0.0999 

400 

17-93 

0.0997 

—  O.OO02 

O.IOOI 

0.0999 

400 

22.92 

0.0997 

—  O.OOO2 

O.IOOI 

0.0999 

400 

18 

o  .  099*7 

—  O.OOO2 

O.  IOOI 

0.0999 

400 

16 

o  .  0996 

—  o  .  0003 

o.  1498 

0.1495 

600 

i 

23.26 

0.1493 

—  O.OOO2 

o.  1498 

0.1495 

600 

i 

16.66 

0.1493 

—  O.OOO2 

0.1498 

0.1495 

ooo 

20.  b; 

O.U75 

—  O.OC2O 

0.1996 

o.  1992 

800 

2 

22.38 

0.  IQQO 

—  O.OOO2 

0.1996 

o  1992 

800 

2 

17.29 

o.  1909 

+  0.0007 

o  1996 

0.1992 

800 

2 

22    2O 

o.  iggr 

—  O.OOOI 

0.4045 

0.4037 

1600 

4 

I6.O3 

0.404.2 

+o  .  0005 

0.4045 

0.4037 

1600 

4 

16.2 

0.4023 

—  0.0014 

0.4018 

0.4010 

1600 

4 

16.24 

O.4007 

-0.0003 

0-5051 

0.5041 

1800 

4 

I5-27 

O.5O26 

—0.0015 

In  applying  the  method  to  an  insoluble  ferric  compound,  like 
the  oxide,  the  material,  best  not  exceeding  the  equivalent  of  0.2 
grm.  of  the  oxide,  is  dissolved  in  hydrochloric  acid;  the  solution  is 


IRON  495 

evaporated  to  pasty  condition  and  then  diluted ;  a  drop  of  potas- 
sium sulphocyanate  is  introduced  with  n/io  sodium  thiosulphate, 
15  cm.3  to  35  cm.3  in  excess;  the  liquid  is  allowed  to  stand  until 
perfectly  colorless;  and  the  excess  of  thiosulphate  is  determined 
by  n/io  iodine  with  the  aid  of  the  starch  indicator. 

Results  obtained  by  this  procedure  are  given  in  the  table. 

The  Standardization  of  Potassium  Permanganate  in  Iron  Analysis. 

Charlotte  F.  Roberts  *  has  discussed  the  standardization  of 
potassium  permanganate  for  practical  work  in  iron  analysis. 
The  best  authorities  agree  in  considering  that  the  standard  of  the 
permanganate  should  finally  be  referred  to  a  standard  solution  of 
ferric  chloride,  but  the  difficulty  consists  in  determining  with 
accuracy  the  amount  of  iron  in  this  solution.  Though  the  purest 
iron  which  can  be  obtained  commercially  is  used  as  the  basis,  the 
resulting  ferric  chloride  still  contains  some  silica  and  phosphorus, 
which  must  be  eliminated  or  the  amount  determined  gravimetri- 
cally.  The  process  of  determining  the  amount  of  iron  in  the 
ferric  chloride  solution,  upon  which  the  potassium  permanganate 
is  finally  standardized,  thus  becomes  long  and  tedious. f  To  obvi- 
ate this,  the  potassium  permanganate  may  be  compared  with  a 
solution  containing  a  known  weight  of  iron  and  the  solution 
of  ferric  chloride  may  then  be  standardized  by  reduction  and 
titration  with  this  same  potassium  permanganate.  Since  electro- 
lytic iron  is  undoubtedly  the  purest  form  of  iron  known,  it  would 
seem  that  potassium  permanganate  titrated  against  this  might 
be  expected  to  give  trustworthy  results  for  the  first  comparison. 

For  the  production  of  a  definite  amount  of  electrolytic  iron  two 
different  courses  are  open.  Either  a  weighed  amount  of  a  pure 
iron  salt,  as  ammonio-ferrous  sulphate,  may  be  taken  and  the 
iron  completely  precipitated  by  electrolysis;  or  an  indefinite 
quantity  of  the  salt  may  be  taken  and  subjected  to  electrolysis 
for  a  time,  and  the  amount  of  iron  determined  by  weighing  the 
electrolytic  deposit.  This  second  method  is  recommended  as 
being  much  more  rapid  and  free  from  details  of  manipulation 
which  render  the  first  difficult.  About  10  grm.  of  ammo.nio- 
ferrous  sulphate  are  dissolved  in  150  cm.3  of  water,  5  cm.3  of  a 

*  Charlotte  F.  Roberts,  Am.  Jour.  Sci.,  [3],  xlviii,  290  (1894). 
t  Blair,  The  Chemical  Analysis  of  Iron,  yth  Ed.,  page  235. 


496 


METHODS  IN  CHEMICAL  ANALYSIS 


saturated  solution  of  potassium  oxalate  are  added,  and  the  whole 
is  heated  with  a  considerable  quantity  of  solid  ammonium  oxa- 
late until  a  clear  solution  is  obtained.  This  solution  is  decom- 
posed in  a  beaker  between  two  platinum  electrodes,  the  iron  being 
deposited  on  a  piece  of  platinum  foil  of  a  size  convenient  for 
insertion  in  a  rather  large  weighing  bottle,  in  which  it  is  weighed 
both  before  and  after  the  electrolysis.  An  hour  and  a  half,  with 
a  current  of  2  amperes,  is  sufficient  for  the  precipitation  of  0.4 
grm.  to  0.5  grm.  of  pure  iron,  and  it  was  found  unadvisable  to  use 
a  current  much  stronger  than  2  amperes,  since  a  higher  current 
showed  a  tendency  to  render  the  deposit  less  smooth  and  com- 
pact. After  washing,  drying  and  weighing  in  the  usual  way, 
the  iron  was  dissolved  in  hydrochloric  acid,  the  weighing  bottle 
being  used  instead  of  the  small  flask  ordinarily  employed  in  this 
operation,  the  oxidized  iron  was  reduced  with  zinc,  and  finally 
titrated  with  the  solution  of  potassium  permanganate  in  presence 
of  sulphuric  acid  and  a  large  amount  of  water. 

The  following  table  shows  the  results  obtained  by  this  proce- 
dure, the  first  column  giving  the  weight  of  the  electrolytic  deposit 
of  iron,  and  the  second  the  weight  of  iron  found  by  titration  with 
potassium  permanganate  previously  standardized  on  a  special 
ammonium  oxalate  which  had  been  shown  to  give  results  iden- 
tical with  those  obtained  by  the  use  of  specially  prepared  anhy- 
drous lead  oxalate: 

Comparison  of  Standards. 


I. 

II. 

Difference. 

Average. 

0-4357 

0.4364 

+O.OO07 

0-3551 

0.3559 

+o  .  0008 

0-2552 

0.2550 

—  0.0002 

o  .  2898 

0.2890 

—  O.OOOS 

0.1590 

0.1599 

+  O.OOO9 

0.3528 

0-3534 

+o  .  0006        f> 

o.oooo 

o  .  4498 

o  .  4494 

—  O.0004 

o  .  5086 

o  .  5085 

—  o.oooi 

0.4462 

0-4457 

—0.0005 

0.4226 

0.4222 

—  0.0004 

0.5170 

0-5165 

—0.0005 

The  standard  of  potassium  permanganate  as  determined  from 
pure  iron  is  identical  with  that  obtained  in  these  experiments 
with  the  special  ammonium  oxalate,  but  the  standard  obtained 


IRON  497 

in  the  former  way  would  under  ordinary  conditions  be  more  satis- 
factory for  work  in  iron  analysis.  A  simple  and  rapid  method, 
then,  for  standardizing  the  potassium  permanganate  solution  is 
to  determine  its  strength,  first,  by  comparison  with  electrolytic 
iron  in  the  manner  above  described.  Then  by  reduction  and 
titration  with  the  permanganate  the  exact  amount  of  iron  in  each 
cubic  centimeter  of  the  ferric  chloride  solution  may  be  determined. 
This  being  ascertained,  the  solution  of  ferric  chloride  may  be 
employed  at  any  time  for  the  standardization  of  potassium 
permanganate. 

The  Behavior  of  Ferric  Chloride  in  the  Jones  Reductor. 

The  column  of  amalgamated  zinc  as  applied  in  the  Jones 
reductor  *  has  proved  very  effective  in  the  reduction  of  ferric 
sulphate  preparatory  to  the  estimation  of  the  ferrous  salt  by 
'potassium  permanganate.  The  impression  has  prevailed  that 
the  salt  of  iron  acted  upon  by  the  amalgamated  zinc  must  be  the 
sulphate  and  that  chlorides  and  nitrates  must  not  be  present  even 
in  small  amounts.  Randall  |  has  shown,  however,  that  it  is 
possible  to  reduce  ferric  chloride  in  the  zinc  reductor  and  to 
determine  the  iron  with  success  by  potassium  permanganate, 
provided  the  titration  is  carried  on  in  the  presence  of  manganous 
sulphate  and  in  solutions  sufficiently  dilute.  A  small  excess  of 
hydrochloric  acid  has  no  influence  on  the  result  in  dilute  solu- 
tions, and  in  a  volume  of  one  liter  the  excess  may  amount  to  as 
much  as  25  cm.3  of  the  strongest  acid. 

According  to  the  method  described,  the  procedure  is  to  first 
run  100  cm.3  of  warm  dilute  2.5  per  cent  sulphuric  acid  through 
the  reductor  charged  with  amalgamated  2O-mesh  zinc,  next  to 
pass  in  the  iron  solution  diluted  with  100  cm.3  of  warm  2.5  per 
cent  sulphuric  acid  and  then  to  wash  down  with  200  cm.3  of  the 
warm  dilute  acid  followed  by  100  cm.3  of  hot  water.  The 
receiving  flask  of  the  reductor  is  kept  in  a  vessel  containing  run- 
ning tap  water,  so  that  the  solution  is  cooled  as  fast  as  it  is 
reduced.  Manganous  sulphate,  I  grm.  to  5  grm.,  is  added  before 
the  titration  with  permanganate.  Following  are  results  obtained 
by  this  method,  and,  for  comparison,  the  results  obtained  upon 
the  same  solution  of  ferric  chloride  standardized  by  evaporating 

*  See  page  437. 

t  D.  L.  Randall,  Am.  Jour.  Sci.,  [4],  xxi,  128. 


498 


METHODS  IN  CHEMICAL  ANALYSIS 


measured  amounts  with  10  cm.3  of  sulphuric  acid  to  the  fuming 
point  of  the  acid,  passing  the  solution  through  the  reductor  and 
titrating  with  permanganate. 

Reduction  After  Conversion  of  Chloride  to  Sulphate. 


Fed,. 

H2S04[i:i.] 

Volume  at 
titration. 

KMnO4. 

Fe  found. 

Variation 
from  average. 

cm.3 

cm.* 

cm.s 

cm.8 

grin. 

grm. 

75 
75 
75 

25 
25 
25 

750 
750 
750 

70.75 
70.83 
70.83 

0.4867] 
0-4873   I 
0.4873   f 

0.4872* 

r  -0.0005 

J    +O.OOOI 
]   +O.OOOI 

75 

25 

75° 

70.88 

0.4876; 

(.+0.0004 

Reduction  of  the  Chloride. 


Fed,. 
cm.3 

Fe  taken.* 
grm. 

H2S04 
2.5  per 
cent. 

cm.3 

HC1 
cone. 

cm.8 

Volume  at 
titration. 

cm.8 

MnSO4. 
grm. 

KMnO4. 
cm.8 

Fe  found, 
grm. 

£*rror. 
grm. 

75 

0.4872 

IOO 

o 

750 

70.81 

0.4871 

—  O.OOOI 

75 

0.4872 

IOO 

o 

750 

70.75 

0.4867 

-0.0005 

75 

0.4872 

IOO 

o 

750 

70.83 

0.4873 

+O.OOOI 

75 

0.4872 

IOO 

0 

750 

70.82 

0.4872 

o.oooo 

75 

0.4872 

IOO 

o 

750 

70.83 

0.4873 

+O.OOOI 

IOO 

0.6497 

IOO 

o 

750 

94-43 

0.6497 

o.oooo 

IOO 

0.6497 

.     IOO 

o 

750 

94-44 

o  .  6498 

+0.0001 

IOO 

0-6497 

IOO 

IO 

750 

94-44 

0.6498 

+O.OOOI 

IOO 

0.6497 

IOO 

20 

1000 

94-53 

0.6503 

+o  .  0006 

IOO 

0.6497 

IOO 

25 

IOOO 

94-53 

o  .  6503 

+o  .  0006 

IOO 

0.6497 

IOO 

25 

IOOO 

•25 

94-53 

o  .  6503 

+0.0006 

IOO 

0.6497 

IOO 

25 

IOOO 

•25 

94.48 

0.6500 

+0.0003 

IOO 

0.6497 

IOO 

25 

IOOO 

5-oo 

94-49 

0.6501 

+o  .  0004 

*  Mean  of  results  obtained  by  titration  after  the  conversion  of  the  chloride  to  sulphate. 

The  Effect  of  Nitric  Acid  in  the  Titration  of  a  Ferrous  Salt  by 
Potassium  Permanganate. 

In  Schneider's  *  method  for  the  determination  of  manganese, 
permanganic  acid  is  titrated  with  hydrogen  peroxide  in  the  pres- 
ence of  nitric  acid,  and  Ibbotson  f  and  Brearley  in  their  modifi- 
cation of  this  process  recommend  the  use  of  standard  ferrous 
ammonium  sulphate  instead  of  hydrogen  peroxide.  Blair  t 
recommends  the  use  of  ferrous  ammonium  sulphate.  Obviously 
it  is  of  interest  to  know  the  exact  effect  of  nitric  acid  upon  solu- 

*  Ding.  Pol.  Jour.f  cclxix,  224. 

f  Chem.  News,  Ixxxiv,  247. 

|  Blair,  Jour.  Am.  Chem.  Soc.,  xxvi,  793.     Also,  Chemical  Analysis  of  Iron. 


IRON 


499 


tions  of  the  ferrous  salt  undergoing  oxidation  by  permanganate, 
and  an  investigation  of  this  matter  by  Randall  *  has  shown  that 
when  more  than  10  per  cent  by  volume  (20  cm.3  in  200  cm.3)  of 
the  concentrated  acid  is  present  oxidation  of  the  ferrous  salt 
takes  place,  as  is  made  plainly  evident  by  the  change  of  color 
of  the  solution,  low  results,  and  uncertain  end  reaction.  The 
error  due  to  oxidation  of  the  ferrous  salt  by  nitric  acid  is,  how- 
ever, in  part  counterbalanced  by  the  reoxidation  by  perman- 
ganate of  any  nitrous  acid  which  may  have  been  produced.  If 
the  titration  is  made  without  unnecessary  delay  the  presence 
of  as  much  as  5  per  cent  by  volume  of  nitric  acid  has  no  appreci- 
able effect  upon  the  estimation  of  ferrous  iron.  Following  are 
results  of  the  titration  of  two  different  solutions  of  ferrous  sul- 
phate in  presence  of  varying  amounts  of  nitric  acid : 

Permanganate  Titration  in  Presence  of  Nitric  Acid. 


FeSO4. 
cm.3 

Dilution. 
cm.3 

H2S04  [i  :  i.] 
cm.3 

HNO3. 
cm.3 

Approx. 
n/io  KMn04. 

cm.3 

25 

2OO 

5 

0 

13.37 

25 

2OO 

5 

o 

13-39 

25 

2OO 

5 

o 

I3.4I 

25 

2OO 

5 

o 

13.38 

T 

* 

25 

2OO 

5 

o 

13.38 

25 

200 

o 

3 

13.38 

25 

200 

o 

5 

13.40 

25 

200 

0 

5 

13.41 

25 

2OO 

0 

5 

13.40 

125 

2OO 

0 

5 

13.38 

fas 

2OO 

5 

o 

13-47 

25 

200 

5 

o 

13.50 

1  25 

200 

5 

o 

13-49 

H^  25 

2OO 

0 

5 

13-53 

25 

2OO 

o 

10 

13.50 

[20 

2OO 

o 

20 

13.51 

200 

o 

30 

13  oo 

The  Permanganate  Estimation  of  Iron  in  Presence  of  Titanium. 

For  analytical  purposes,  a  ferric  salt  in  solution  is  most  easily 
and  conveniently  reduced  to  the  ferrous  condition  by  the  action 
of  zinc ;  and  where  many  determinations  of  iron  are  to  be  made, 
the  use  of  the  well-known  Jones  reductorf  yields  accurate  results 

*  D.  L.  Randall,  Am.  Jour.  Sci.,  [4],  xxiii,  139. 
t  See  page  347. 


500  METHODS  IN  CHEMICAL  ANALYSIS 

very  rapidly.  The  use  of  zinc,  whether  in  the  flask  or  in  the 
redactor,  has,  however,  been  precluded  when  the  ferric  salt  is 
accompanied  by  titanic  acid,  for  this  substance  is  reduced  with 
the  iron  and  subsequently  oxidized  by  the  permanganate  in  the 
titration  process.  When,  therefore,  titanium  is  present  with  the 
iron,  it  has  been  customary  to  have  recourse  to  other  methods  of 
reduction.  In  this  event,  either  hydrogen  sulphide  or  sulphur 
dioxide  is  substituted  for  the  zinc  to  bring  about  the  reduction 
of  the  ferric  salt,  since  titanic  acid  is  not  reduced  by  these  re- 
agents ;  but  the  removal  of  the  excess  of  hydrogen  sulphide  or  of 
sulphur  dioxide  from  solution  without  oxidation  of  the  ferrous 
salt  is  not  an  easy  or  rapid  process. 

Gooch  and  Newton  *  have  studied  the  problem  of  adapting  the 
ordinary  convenient  process  of  reducing  the  ferric  salt  by  zinc 
to  the  estimation  of  iron  in  presence  of  titanium.  It  is  obvious 
that  to  solve  this  problem  it  is  only  necessary  to  find  and  employ 
some  reagent  which  shall  be  neutral  toward  the  ferrous  salt  but 
capable  of  reoxidizing  the  titanium  compounds  formed  by  the 
reducing  action  of  the  zinc,  and  shall  have  no  action  on  the  per- 
manganate. Compounds  of  silver,  copper  or  bismuth  oxidize 
easily  the  reduced  titanium  salt;  but  the  use  of  a  compound  of 
silver  is  precluded  by  the  fact  that  it  oxidizes  also  the  ferrous 
salt  to  some  extent  as  well  as  the  titanium  salt.  Cupric  salts 
and  pure  bismuth  oxide  prove,  however,  to  be  without  action 
upon  the  ferrous  salt.  It  is  found  that  the  violet  color  of  the  solu- 
tion containing  the  titanium  compound  produced  by  the  action 
of  zinc  upon  the  titanium  sulphate  is  discharged  upon  adding  a 
little  cupric  sulphate  to  the  solution  and  heating,  and,  after 
filtering,  a  drop  of  potassium  permanganate  gives  its  character- 
istic rose  tint  to  the  solution.  It  is  found  also  that  when  cupric 
oxide  is  added  to  a  similarly  reduced  solution  of  the  titanium  salt 
the  characteristic  color  vanishes  on  shaking  the  solution.  The 
following  table  contains  the  results  obtained  in  titrating  with 
potassium  permanganate  the  ferrous  salt  left  after  reducing  by 
zinc  in  small  flasks  carefully  measured  amounts  of  ferric  sulphate 
and  titanium  sulphate,  treating  the  mixture  thus  obtained  with 
cupric  sulphate  or  with  cupric  oxide,  and  filtering  off  the  reduced 
copper  and  cuprous  salt: 

*  F.  A.  Gooch  and  H.  D.  Newton,  Am.  Jour.  Sci.,  [4],  xxiii,  365.    . 


IRON 


501 


Oxidation  of  the  Titanous  Sulphate  by  Cupric  Compounds:    Titration  of  Ferrous 
Sulphate  by  Permanganate. 


Fe2O3  taken, 
grm. 

TiO2  taken, 
grm. 

Fe2O3  found, 
grm. 

Error, 
grm. 

O.I37S 
0.1375 
0.1375 

O.  I 
O.I 
O.I 

0.1378 

0.1374 
0.1377 

+0.0003  } 
—  O.OOOI  > 
+0.0002  ) 

Treated  with 
CuSO4. 

0.1375 
0-1375 
0.1375 

O.  I 
O.I 
O.  2 

0.1378 
0.1378 
0.1382 

+0.0003  ) 
+0.0003  > 
+0.0007  ) 

Treated  with 
CuO. 

Similar  experiments  in  which  bismuth  oxide  was  substituted 
for  the  copper  oxide  are  also  recorded.  To  the  measured  amount 
of  ferric  sulphate  and  titanium  sulphate  contained  in  a  small 
flask,  provided  as  usual  with  the  funnel  valve,  zinc  is  added  and 
the  reduction  effected  in  the  ordinary  manner.  The  titanium 
salt  appears  to  act  catalytically  in  this  process,  so  that  reduction 
goes  on  more  easily  and  with  less  expenditure  of  zinc  than  in  the 
similar  reduction  of  the  ferric  salt  taken  by  itself.  After  the  zinc 
has  disappeared,  the  solution,  of  characteristic  violet  color,  is 
cooled  in  the  flask,  treated  with  a  little  bismuth  oxide,  gently 
shaken,  filtered  from  the  excess  of  bismuth  oxide  and  the  precipi- 
tated bismuth  into  about  a  liter  of  cold  water,  and  titrated  with 
standard  potassium  permanganate.  Results  are  given  in  the 
following  table: 

Oxidation  of  Titanous  Sulphate  by  Bismuth  Oxide:   Titration  of  Ferrous  Sulphate 

by  Permanganate. 


Ferric 
sulphate. 

cm.8 

TiO2. 

grm. 

KMn04. 
cm.a 

Fe-jOa  taken. 

grrn. 

Fe2O3  found, 
grm. 

Error, 
grm. 

IO 

0.04 

12.84 

0.0993 

0.0992 

—  O.OOOI 

IO 

0.06 

12.85 

0.0993 

0.0993 

O.OOOO 

10 

0.08 

12.90 

0.0993 

0.0997 

+0.0004 

10 

O. 

12.90 

0.0993 

0.0997 

+0.0004 

10 

O. 

12.89 

0.0993 

o  .  0996 

+O.OOO3 

IO 

0. 

12.85 

0.0993 

0.0993 

O.OOOO 

IO 

0. 

12.80 

0.0993 

o  .  0989 

—  0.0004 

IO 

O. 

12.90 

0.0993 

0.0997 

+0.0004 

10 

O.2 

12.90 

0.0993 

0.0997 

+0.0004 

10 

O.  2 

12.89 

o  .  0993 

0.0996 

+0.0003 

10 

0.2 

12.90 

0.0993 

0.0997 

+0.0004 

20 

O.  I 

25.70 

0.1986 

0.1986 

O.OOOO 

502 


METHODS  IN  CHEMICAL  ANALYSIS 


It  is  evident  that  either  cupric  sulphate,  cupric  oxide,  or 
bismuth  oxide  may  be  used  to  reoxidize  the  salt  of  titanium 
reduced  by  zinc,  without  affecting  appreciably  the  ferrous  salt 
in  solution. 

Similar  results  are  obtained  when  the  ferric  sulphate  solution 
is  passed  through  the  column  of  amalgamated  zinc  in  the  Jones 
reductor.*  The  flask  is  kept  cool  in  running  water,  a  small 
amount  of  bismuth  oxide  added,  the  flask  shaken  and  allowed  to 
stand  a  few  minutes,  and  the  mixture  filtered  with  the  aid  of  the 
suction  pump.  In  the  cold  solution  free  from  dissolved  oxygen 
there  is  little  danger  of  reoxidation  of  ferrous  sulphate,  as  has  been 
shown  by  Peters  and  Moody. f  The  experimental  results  are  given 
in  the  table. 

The  Zinc  Reductor:  Bismuth  Oxide:   Permanganate. 


Ferric 
sulphate. 

cm.» 

TiOj. 

grm. 

KMn04. 
cm.1 

FezOa  taken, 
grin. 

Fe,O,  found, 
grm. 

Error, 
grm. 

40 

O.OI 

46.80 

0.3943 

0-3943 

O.OOOO 

40 

O.O2 

46.79 

0.3943 

0.3942 

—  O.OOOI 

40 

0.04 

46.80 

0.3943 

0-3943 

0.0000 

40 

0.06 

46.83 

0.3943 

0.3946 

+0.0003 

40 

o. 

46.75 

0.3943 

0-3939 

—0.0004 

40 

o. 

46.82 

0.3943 

0-3945 

+O.OOO2 

40 

0. 

46.78 

0.3943 

0.3941 

—  O.OOO2 

40 

0. 

46.80 

0.3943 

0-3943 

0.0000 

40 

o. 

46.75 

0.3943 

0-3939 

—  0.0004 

40 

o. 

46.80 

0-3943 

0-3943 

O.OOOO 

40 

0.2 

46.77 

0-3943 

0.3940 

—0.0003 

40 

0.2 

46.81 

0-3943 

0-3944 

+0.0001 

40 

O.2 

46.85 

0-3943 

0-3947 

+0.0004 

The  Estimation  of  Iron  by  Potassium  Permanganate  after  Reduc- 
tion with  Titanous  Sulphate. 

Knecht  \  was  the  first  to  recommend  the  use  of  titanium 
sesquioxide  and  its  salts  in  volumetric  operations  where  a  rapid 
and  powerful  reducing  agent  is  required,  and  later  in  collabo- 
ration with  E.  Hibbert  §  published  a  method  for  the  direct  titra- 
tion  of  ferric  chloride  by  a  standard  solution  of  titanous  chloride, 

*  See  page  347. 

t  See  page  371. 

J  Ber.  Dtsch.  Chem.  Ges.  xxvi,  166. 

§  Ibid.,  xl,  3819. 


IRON  503 

using  potassium  sulphocyanate  as  an  indicator,  the  reaction 
between  the  two  salts  taking  place  according  to  the  following 
equation : 

FeCl3  +  TiCl3  =  FeCl2  +  TiCl4. 

According  to  these  investigators  the  method  yields  excellent 
results,  and  rapidly.  The  only  precautions  necessary  are  that 
the  solution  of  titanous  chloride,  being  naturally  very  sensitive 
to  the  action  of  atmospheric  oxygen,  must,  after  having  been 
boiled  with  marble  to  expel  occluded  oxygen,  be  kept  under  a 
constant  pressure  of  hydrogen.  It  has  been  found,  however,  that 
even  with  such  precautions  the  standard  of  the  solution  grad- 
ually changes  and  must  be  checked  from  time  to  time  against 
known  amounts  of  ferric  iron.  The  proposal  has  therefore  been 
made  by  Newton,*  to  reduce  the  iron  by  titanous  sulphate, 
oxidize  the  excess  of  titanous  sulphate,  and  titrate  the  remaining 
ferrous  salt  by  permanganate. f 

A  solution  of  titanous  sulphate  of  convenient  strength  may 
be  made  up  by  mixing  20  grm.  of  commercial  titanic  acid  with 
three  times  its  own  weight  of  a  mixture  of  sodium  and  potassium 
carbonates  and  fusing  in  a  platinum  crucible,  treating  the  melt 
(after  being  finely  ground)  in  a  platinum  dish  with  hot  concen- 
trated sulphuric  acid,  cooling,  diluting  a  little,  filtering  through 
asbestos,  treating  with  zinc  until  reduction  is  accomplished,  and, 
while  zinc  is  still  left  in  the  flask,  filtering  the  solution  quickly 
through  a  platinum  cone  into  about  two  liters  of  freshly  distilled 
water  contained  in  a  small  reservoir  connected  with  burette  and 
hydrogen  generator. 

To  determine  ferric  sulphate  in  solution  it  is  only  necessary  to 
add  in  the  cold  an  excess  of  titanous  sulphate,  prepared  as  de- 
scribed, destroy  this  excess  by  treating  with  a  little  bismuth 
oxide,  t  filter  from  the  excess  of  bismuth  oxide  into  about  a  liter 
of  cool  distilled  water,  and  titrate  with  permanganate.  If  an 
appreciable  amount  of  hydrochloric  acid  is  present  in  the  solution 
of  the  ferric  salt  it  is  advisable  to  evaporate  to  dryness  and  con- 
vert the  chloride  to  sulphate  by  use  of  concentrated  sulphuric 
acid.  Upon  addition  of  the  sulphuric  acid  a  white  pasty  mass 
is  formed  which  rapidly  goes  into  solution  on  diluting  with  water 

*  H.  D.  Newton,  Am.  Jour.  Sci.,  [4],  xxv,  343. 
t  See  page  500. 
t  See  page  501. 


504 


METHODS  IN  CHEMICAL  ANALYSIS 


and  warming.  As  titanous  sulphate  prepared  in  the  manner 
described  above  always  contains  some  iron,  it  is  necessary  to 
make  a  correction  for  this  by  treating  with  bismuth  oxide 
an  amount  of  the  solution  equal  to  that  used,  filtering,  and 
running  in  potassium  permanganate  to  color.  This  correction 
should  not  amount  to  more  than  o.i  cm.3  when  working  with 
0.3  grm.  of  ferric  oxide.  If  these  simple  precautions  be  taken, 
ferric  iron  may  be  determined  with  rapidity  and  exactness  by 
reduction  with  titanous  sulphate,  treatment  with  bismuth  oxide, 
and  titration  with  potassium  permanganate.  Test  results  are 
given  in  the  table. 

Permanganate  Titration  after  Reduction  with  Titanous  Sulphate. 


Fe2(S04)3. 
cm.* 

KMnO4. 
cm  .s 

Fe2O3  taken. 

grm. 

Fe2O3  found, 
grm. 

Error, 
grm. 

2O 

I3-46 

0.1063 

o.  1064 

+O.OOOI 

2O 

I3-42 

0.1063 

o.  1060 

—  0.0003 

20 

13-44 

0.1063 

O.IO62 

—  0.0001 

20 

13-44 

0.1063 

O.IO62 

—  O.OOOI 

20 

I3-4I 

o  .  1063 

o.  1060 

-0.0003 

30 

20.  l8 

0.1594 

0.1594 

o.oooo 

30 

2O.2O 

0-1594 

0.1596 

+O.OOO2 

30 

2O.  2O 

0.1594 

0.1596 

+0.0002 

30 

2O.22 

0.1594 

0.1598 

+0.0004 

30 

2O.  IQ 

0.1594 

0.1595 

+O.OOOI 

40 

26.92 

0.2127 

O.2I27 

O.OOOO 

40 

26.90 

0.2127 

9.2125 

—  O.OO02 

40 

26.90 

0.2127 

0.2125 

—  0.0002 

40 

26.93 

0.2127 

0.2128 

-fo.oooi 

40 

26.90 

0.2127 

O.2I25 

—  O.OO02 

Separations  of  Iron  by  Volatilization  in  Gaseous  Hydrogen 

Chloride. 

Metallic  iron  is  easily  acted  upon  by  an  excess  of  chlorine  at 
moderately  elevated  temperatures  with  the  formation  of  ferric 
chloride  and  by  hydrochloric  acid  gas  with  formation  of  ferrous 
chloride.  Out  of  contact  with  air,  or  moisture,  both  chlorides 
may  be  volatilized  at  appropriate  temperatures  —  the  ferric  chlo- 
ride below  200°  C. ;  the  ferrous  chloride  at  a  bright  red  heat. 
If  water  vapor,  or  oxygen,  or  air,  be  present  during  the  heating, 
both  chlorides  are  partially  decomposed  with  the  formation  of 
non-volatile  residues,  ferric  oxide  or  ferric  oxychloride.  When 


IRON  505 

ferric  oxide  is  submitted  to  the  action  of  hydrochloric  acid  gas 
at  about  200°  the  greater  part  of  the  iron  sublimes,*  as  ferric 
chloride,  but  a  residue  remains.  At  the  outset  the  ferric  oxide 
volatilizes  quickly  and  abundantly  in  the  form  of  the  greenish 
vapor  of  ferric  chloride,  and  if  the  operation  is  interrupted  at  this 
stage  the  residue  which  remains  is  nearly  black,  insoluble  in 
water,  slightly  soluble  in  cold  hydrochloric  acid,  and  readily 
soluble  in  hot  hydrochloric  acid  with  the  formation  of  ferric 
chloride.  It  is  probably  something  analogous  to  the  oxychlo- 
ride  which  Rousseau  f  identified  as  the  product  of  the  action  of 
water  upon  ferric  chloride  at  275°  to  300°.  This  dark  residue 
yields  to  the  action  of  the  hydrochloric  acid  at  180°  to  200°  only 
slowly;  but  ultimately  the  residue  is  essentially  ferrous  chloride. 
Little  volatilization  occurs  within  the  range  of  temperature  from 
200°  to  500°. 

If  the  temperature  of  the  oxide  is  450°  to  500°  when  the  brisk 
current  of  acid  begins  to  act,  the  whole  mass  of  oxide  is  rapidly 
converted  and  volatilizes  without  residue,  the  production  of  the 
ferrous  chloride  (formed  by  dissociation  of  ferric  chloride)  being 
apparently  kept  at  a  minimum  by  the  adjustment  of  equilibrium 
in  the  atmosphere  of  ferric  chloride  and  chlorine  resulting  from 
the  partial  dissociation.  If  dissociation  of  ferric  chloride  to 
ferrous  chloride  is  the  cause  of  the  formation  of  a  residue  at  200°, 
the  temperature  of  slow  action,  the  introduction  of  chlorine  into 
the  atmosphere  of  hydrochloric  acid  gas  should  change  the  condi- 
tion of  equilibrium  and  enable  the  ferric  chloride  to  volatilize 
without  dissociation.  Gooch  and  Havens  J  find,  as  a  matter  of 
fact,  that  if  a  little  manganese  dioxide  is  added  to  the  contents  of 
the  hydrogen  chloride  generator,  so  that  the  gas  may  carry  with 
it  a  little  chlorine,  every  trace  of  ferric  oxide  is  volatilized  from 
the  boat  at  180°  to  200°.  The  residue  of  ferrous  chloride  found 
at  1 80°  to  200°  when  the  hydrochloric  acid  is  used  alone  is  like- 
wise volatilized  at  the  same  temperature  when  the  admixture 
of  chlorine  is  made. 

These  facts,  that  ferric  oxide  is  completely  volatile  in  hydro- 
chloric acid  gas  applied  at  once  at  a  temperature  of  450°  to  500°, 
and  at  180°  to  200°  if  the  acid  carries  a  little  chlorine,  open 

*  Moyer,  Jour.  Am.  Chem.  Soc.,  xviii,  1029. 

t  Compt.  rend.,  cxvi,  118. 

J  F.  A.  Gooch  and  Franke  Stuart  Havens,  Am.  Jour.  Sci.,  [4],  vii,  370. 


506  METHODS  IN  CHEMICAL  ANALYSIS 

the  way  to  many  analytical  separations  of  iron  from  substances 
not  volatile  under  these  conditions. 

The  separation  of  the  iron  oxide  from  various  oxides  proves  to 
be  complete  at  450°  to  500°  if  the  mixture  is  submitted  at  once  to 
the  action  of  hydrochloric  acid  gas,  or  at  180°  to  200°  when 
chlorine  is  mixed  with  the  hydrochloric  acid.  The  temperature 
of  red  heat  employed  by  Deville  *  is  unnecessary  if  the  mixed 
oxides  are  submitted  at  once  to  the  action  of  hydrochloric  acid 
at  450°  to  500°  without  previous  gentle  heating  in  the  acid 
atmosphere.  The  mixture  of  chlorine  and  hydrochloric  acid  is 
to  be  preferred,  however,  not  only  because  the  temperature  of 
the  reaction  is  lower  but  because  it  needs  no  regulation,  while 
the  danger  of  error  arising  from  the  liability  of  ferric  chloride 
to  dissociate,  or  from  deficiency  of  oxidation  in  the  oxide 
treated,  or  from  mechanical  loss  due  to  too  rapid  volatilization 
is  avoided. 

According  to  the  procedure  adopted,  the  mixed  oxides,  put  in 
a  porcelain  boat  which  is  placed  in  a  wide  combustion  tube  heated 
in  a  small  furnace,  are  submitted  to  the  action  of  dry  hydrochloric 
acid  gas,  generated  by  dropping  sulphuric  acid  upon  a  mixture  of 
strong  hydrochloric  acid,  common  salt,  and  a  small  amount  of 
manganese  dioxide.  -  The  gas  is  admitted  at  one  end  of  the  com- 
bustion tube  and  passed  out  at  the  other  through  a  water  trap, 
while  the  required  temperature,  best  200°  to  300°,  is  maintained 
by  regulating  the  burners  of  the  furnace.  The  time  of  action 
varies  somewhat  with  the  condition  of  the  oxide  to  be  volatilized 
and  the  temperature.  Generally  an  hour's  heating  at  200° 
proves  sufficient  for  the  complete  removal  of  o.i  grm.  of  iron. 
At  higher  temperatures  the  action  is  more  rapid;  but  the  non- 
volatile oxide  is  liable  to  mechanical  loss  if  the  volatilization  of 
the  iron  is  too  rapid.  It  is  better,  therefore,  to  make  use  of  a 
lower  temperature  until  the  volatilization  of  iron  is  nearly  com- 
plete, and  then  to  raise  the  heat  for  a  few  minutes  to  insure  the 
removal  of  the  last  traces  of  the  volatile  chloride. 

Results  of  this  procedure  follow. 

iron  and  Separations  of  iron   oxide  and  aluminium  oxide 

Aluminium.        j^y  fae  procedure  outlined  above  are  shown  in  the 
following  table  :f 

*  Ann.  Chim.,  [3],  xxxviii,  23. 
t  Gooch  and  Havens,  loc.  cit. 


IRON 

Volatilization  of  Ferric  Chloride. 


507 


Fe203 
taken. 

grm. 

A1203 
taken. 

grm. 

A1203 
found. 

grm. 

Error, 
grm. 

Time, 
hours. 

Temperature. 

c°. 

Atmosphere. 

0.1000 
0.2000 
O.IO2O 
0.2145 

O    IOOO 

O.IOI5 
O.  1006 

0.1015 
0.1008 

o.oooo 
o.oooo 
o.oooo 

+0.0002 

o  oooo 

\ 

I 

I 

1 

4 

450-500 
450-500 
450-500 
450-500 

180—200 

HC1. 
HC1. 
HC1. 
HC1. 

HCl+Cla 

O.IOOO 

0.1072 
0.2045 

o.  1050 
o  .  2008 

0.1032 
0.1013 
0.1032 
0.1023 
O.  1007 
0.1087 

0.1032 
o.  1015 

0.1033 
O.IOI9 

o.  1006 

O.IO87 

o  .  oooo 

+O.OOO2 

-j-o.oooi 
—0.0004 

—  O.OOOI 

o.oooo 

I 

ly 

if 

180-200 
180-200 
180-200 

450-500 
450-500 
4<o—  <;oo 

HC1+C12. 
HC1+C12. 
HC1+C12. 
HC1+C12. 
HC1+C12. 
HC1+C12. 

iron  and  The  separation  of  iron  and  beryllium  has  been 

Beryllium.         tested  by  Havens  and  Way.* 

Volatilization  of  Ferric  Chloride. 


Fe2O3  taken, 
grm. 

BeO  taken, 
grm. 

BeO  found, 
grm. 

Error, 
grm. 

0.1309 
0.1285 
o  .  0456 

0.1099 
0.1080 
0.1305 

o.  1081 

0.1311 
0.1285 

0.0457 
0.1099 
0.1081 
0.1290 
0.1083 

4-O.OOO2 
O  .  OOOO 
+O.OOOI 
O.OOOO 
+O.OOOI 

—0.0015 

+0.0002 

0.0997 
o  .  1045 

0.1215 
0.1510 
o  .  2030 

iron  and               Tests  of  the  separation  of  iron  and  chromium  have 
chromium.         aiso  keen  made  by  Havens  and  Way.f 

Volatilization  of  Ferric  Chloride. 

Fe2O3  taken, 
grm. 

Cr2O3  taken, 
grm. 

Cr2O3  found, 
grm. 

Error, 
grm. 

O.IOO8 

O.IOO8 

O  .  OOOO 

o  .  1007 

o.  1006 

0.1006 

O.OOOO 

o  .  1007 

O.IOOO 

O.IOO2 

+O.OOO2 

O.  IOIO 

O.IOO5 

0.1003 

—  0.0002 

o.  1019 

o.  1006 

0.1005 

—  O.OOOI 

0.2007 

0.1003 

0.0999 

—0.0004 

*  Franke  Stuart  Havens  and  Arthur  Fitch  Way,  Am.  Jour.  Sci.,  [4],  viii,  217. 
t  Loc.  cit. 


METHODS  IN  CHEMICAL  ANALYSIS 


Iron  and 
Zirconium. 


The  following  table  contains  results  by  Havens 
and  Way  *  in  the  separation  of  iron  and  zirconium  : 

Volatilization  of  Ferric  Chloride. 


Fe2O3taken. 
grm. 

ZrO2  taken, 
grm. 

ZrOj  found, 
grm. 

Error, 
grm. 

0.1516 

0.1516 

O.OOOO 

O.IOS3 

O.  IOIO 

O.IOIO 

o.oooo 

0.1204 

0.1519 

0.1523 

+0.0004 

0.1236 

0.1516 

0.1517 

-f-O.OOOI 

0.2150 

0.1517 

0.1519 

+  O.OOO2 

The  Estimation  of  Iron  and  Vanadium  in  Presence  of  Each 

Other. 

In  view  of  the  difficulties  attendant  upon  the  separation  of 
iron  and  vanadium,  Edgar  f  has  developed  a  method  by  which 
these  elements  may  be  estimated  in  the  presence  of  each  other. 

If  a  solution  containing  vanadic  acid  and  iron  be  reduced  by 
means  of  sulphur  dioxide  the  reoxidation  by  potassium  perman- 
ganate proceeds  according  to  the  equation 

I.          5  V2O4  +  10  FeO  +  4  KMnO4  =  5  V2O5  +  5  Fe2O3 
+  2  K2O  +  4  MnO. 

If  this  solution,  after  titration,  be  passed  through  a  column  of 
amalgamated  zinc  in  the  Jones  reductor,J  the  receiving  flask  being 
charged  with  a  solution  of  ferric  sulphate,  the  reduction  is  carried, 
in  the  case  of  vanadic  acid,  to  the  condition  of  V2O2  and  the 
reoxidation  by  permanganate  proceeds  according  to  the  equation 

II.          5  V2O2  +  10  FeO  +  8  KMnO4  =  5  V2O5  +  5  Fe2O3 
+  4  K2O  +  8  MnO. 

The  difference  between  the  amounts  of  permanganate  used  in 
the  first  and  second  titrations  is  evidently  used  in  oxidizing  the 
vanadium  from  the  condition  of  V2O2  to  V2O4,  and  measures  the 
amount  of  vanadic  acid  present.  This  being  known,  the  iron 
present  may  be  calculated  from  the  amount  of  permanganate 
used  in  either  titration. 

*  Loc.  cit. 

t  Graham  Edgar,  Am.  Jour.  Sci.,  [4],  xxvi,  79. 

t  See  page  349. 


IRON 


509 


Into  the  slightly  acid  solution  of  the  ferric  salt  and  vanadic  acid, 
contained  in  a  stoppered  flask  provided  with  inlet  and  outlet  tubes, 
is  passed  a  current  of  sulphur  dioxide  until  the  color  changes  from 
red  into  green  and  finally  into  a  clear  blue.  A  few  cubic  centimeters 
of  dilute  sulphuric  acid  are  added,  and  the  solution  is  heated  to 
boiling,  the  current  of  sulphur  dioxide  being  replaced  by  one  of  air- 
free  carbon  dioxide.  When  the  last  traces  of  sulphur  dioxide  have 
been  removed,  the  flask  is  cooled  in  running  water,  the  atmos- 
phere of  carbon  dioxide  being  maintained,  and,  when  thoroughly 
cool,  titrated  with  potassium  permanganate  until  the  color  has 
changed  from  blue  into  yellowish  green.  The  solution  is  then 
heated  to  70°  or  80°  and  the  titration  is  completed  at  that  tem- 
perature. The  amount  of  permanganate  used  indicates  the 
amount  of  vanadic  acid  present  according  to  equation  I. 

Differential  Reductions  by  Sulphur  Dioxide  and  by  Zinc. 


w/io  X  9545- 

IT  f\ 

\7  O 

I. 

II. 

V2OS 
taken. 

found. 

Error  on 
V,06. 

Fe2O, 
taken. 

Fe,0, 

found  . 

Error 
Fe20t. 

KMnO4. 

KMnO4. 

cm.8 

cm.8 

grm. 

grm. 

grm. 

gnu. 

grm. 

grm. 

31.90 

58.02 

0.1136 

0.1137 

+0.0001 

0.1437 

0.1436 

—  O.OOOI* 

31.90 

58.04 

0.1136 

0.1138 

+O.OOO2 

0.1437 

0-J435 

—  O.OOO2* 

31.85 

58.00 

0.1136 

0.1138 

+O  .  OOO2 

0.1437 

0.1433 

—  0.0004* 

31.90 

58.00 

0.1136 

0.1136 

O.OOOO 

0.1437 

0.1437 

O.OOOO* 

25-30 

38.25 

0.0568 

0.0568 

o  .  oooo 

0.1437 

0.1423 

—0.0004' 

25.29 

38.30 

0.0568 

0.0566 

—  O.OOO2 

0.1437 

0.1433 

—0.0004' 

15.98 

29.02 

0.0568 

0.0568 

O.OOOO 

0.0719 

0.0721 

+O.OOO2' 

38.50 

77-60 

0.1704 

0.1702 

—  0.0002 

0.1437 

0.1442 

+0.0005: 

38.45 

77.60 

0.1704 

0.1704 

O.OOOO 

0.1437 

0.1438 

+0.0001; 

38.45 

77.58 

0.1704 

0.1703 

—  O.OOOI 

0.1437 

0.1439 

+0.0002; 

22.50 

48.60 

0.1136 

0.1136 

O.OOOO 

0.0719 

0.0720 

+0.0001* 

22.50 

48.60 

0.1136 

0.1136 

O.OOOO 

0.0719 

0.0720 

+O.OOOI* 

22.45 

48.58 

0.1136 

0.1137 

+0.0001 

0.0719 

0.0716 

—0.0003* 

15-97 

29.07 

0.0568 

0.0570 

+0.0002 

0.0719 

0.0718 

—  O.OOOI  f 

*  35  cm.8  of  a  10  per  cent  solution  of  ferric  alum  in  the  receiver, 
t  20  cm.8  of  a  10  per  cent  solution  of  ferric  alum  in  the  receiver, 
t  50  cm.8  of  a  10  per  cent  solution  of  ferric  alum  in  the  receiver. 

The  solution  just  titrated  for  vanadic  acid,  having  now  a 
volume  of  100  cm.3  to  150  cm.3,  is  passed  through  a  column  of 
amalgamated  zinc  in  a  long  Jones  reductor,  being  preceded  by 
150  cm.3  of  hot  dilute  (2^  per  cent)  sulphuric  acid  and  followed 
by  100  cm.3  of  the  same  acid,  and  finally  200  cm.3  of  distilled 
water.  The  receiving  flask,  containing  an  excess  of  ferric  sul- 


510  METHODS  IN  CHEMICAL  ANALYSIS 

phate,  is  kept  cool  by  means  of  running  water,  and  its  contents, 
.after  the  addition  of  sirupy  phosphoric  acid  to  remove  the  color 
of  the  iron,  are  titrated  with  permanganate  until  the  color  has 
changed  from  bluish  green  to  yellow,  and  the  color  of  the  per- 
manganate begins  to  be  persistent  and  destroyed  only  by  shak- 
ing. The  flask  is  then  heated  to  70°  or  80°  and  the  titration 
completed  in  the  hot  solution. 

The  results  given  in  the  table  show  that  iron  and  vanadium 
may  be  readily  estimated  in  the  presence  of  each  other  by  two 
reductions  —  the  first  with  sulphur  dioxide  and  the  last  with 
amalgamated  zinc,  under  the  conditions  described  above  —  each 
followed  by  an  oxidation  with  permanganate. 

The  Estimation  of  Ferric  Iron,  Vanadic  Acid  and  Chromic  Acid 
in  the  Presence  of  One  Another. 

The  estimation  of  vanadic  acid,  chromic  acid  and  ferric  iron 
associated  with  one  another  has  been  accomplished  by  Edgar  * 
by  making  use  of  processes  of  differential  reduction. 

By  the  action  of  hydrobromic  acid,  vanadic  acid  and  chromic 
acid  are  reduced  according  to  the  equations 

V2O5  +  2  H3r  =  V2O4  +  H2O  +  Br2, 
V2O5  -f  2  CrO3  +  8  HBr  =  V2O4  +  Cr2O3  +  4  H2O  +  4  Br2, 

the  iron  being  unaffected.  By  the  action  of  hydriodic  acid  upon 
the  residue  after  action  of  the  hydrobromic  acid  vanadium 
tetroxide  is  further  reduced  and  the  iron  is  reduced  to  the  ferrous 
condition  while  the  chromic  salt  is  unaffected,  as  shown  in  the 
following  equations: 

V204  +  2  HI  =  V203  +  H20  +  I2, 
Fe203  +  2  HI  =  2  FeO  +  H2O  +  I2. 

If  the  halogen  liberated  in  the  two  reductions  be  separately 
determined,  two  equations  result,  of  which  the  first  gives  the  sum 
of  the  vanadic  and  chromic  acids  and  the  second  the  sum  of  the 
vanadic  acid  and  the  iron ;  the  halogen  equivalent  to  the  vana- 
dium being  the  same  in  each  case.  If  then  either  the  vanadium, 
iron  or  chromium  be  estimated  separately,  a  third  equation  is  ob- 
tained, from  which,  with  the  aid  of  the  first  two,  all  three  con- 
stituents may  be  calculated.  Edgar  found  that  the  estimation 
*  Graham  Edgar,  Am.  Jour.  Sci.,  [4],  xxvii,  174. 


IRON  511 

of  the  chromium  in  a  separate  portion  by  a  modification  of 
Browning's  method  of  reduction  with  arsenious  acid  *  affords 
a  successful  solution  to  this  problem. 

The  distillation  flask  used  by  Edgar  consists  of  a  100  cm.3 
pipette  modified  as  shown  in  Fig.  5,f  the  inlet  tube  being  bent 
upwards  and  having  a  separatory  funnel  sealed  to  its  end,  while 
the  outlet  tube  is  bent  upwards  and  then  down  to  enter  the 
absorption  flask,  a  small  bulb  being  blown  in  it  to  prevent  me- 
chanical loss  during  distillation.  In  carrying  out  the  process  a 
slow  current  of  hydrogen  from  a  Kipp  generator  is  kept  up  through 
the  apparatus,  and  this,  entering  near  the  bottom  of  the  flask, 
makes  it  possible  to  carry  the  distillation  to  very  small  volume 
without  danger  of  "bumping." 

The  entire  process  in  detail  is  as  follows: 

(I)  The  solution,  of  about  50  cm.3  volume,  containing  the  vana- 
date,  chromate  and  ferric  salt,  is  divided  into  two  equal  parts,  one 
of  which  is  placed  in  a  distillation  flask.     To  this,  one  to  two  grams 
of  potassium  bromide  are  added,  together  with  25  cm.3  of  concen- 
trated hydrochloric  acid,  and  the  mixture  is  distilled  until  a  vol- 
ume of  about  25  cm.3  is  reached,  the  reduction  having  always  been 
found  to  be  complete  in  that  time.     The  bromine  liberated  is 
absorbed  in  alkaline  potassium  iodide,  and,  after  cooling  and  acid- 
ifying, the  iodine  liberated  is  estimated  by  titration  with  approx- 
imately tenth  normal  sodium  thiosulphate. 

(II)  The  absorption  apparatus  is  replaced,  and,  after  the  addi- 
tion to  the  solution  in  the  distillation  flask  of  I  grm.  of  potas- 
sium iodide,    10   cm.3  of  concentrated    hydrochloric   acid   and 
3  cm.3  to  5  cm.3  of  sirupy  phosphoric  acid,  distillation  is  again 
carried  on  until  a  volume  of  10  cm.3  has  been  reached;  the  iodine 
thus  liberated  being  estimated  as  before. 

(III)  In  the  second  half  of  the  original  solution  the  chromic 
acid  is  estimated  by  adding  sulphuric  acid  to  slight  acidity, 
3  cm.3  of  sirupy  phosphoric  acid,  and  an  amount  of  standard 
arsenious  acid  in  excess  of  that  required  to  effect  the  reduction  of 
the  chromic  acid.     The  use  of  phosphoric  acid  causes  the  iron  to 
be  precipitated  as  phosphate,  and  thus  the  difficulty  mentioned 
by  Browning  J  in  observing  the  end  point,  due  to  the  reddish- 

*  See  page  407. 
t  See  page  5. 
t  See  pdge  408. 


METHODS  IN  CHEMICAL  ANALYSIS 


i 


=8 


a 


4 

'.£ 
L 


s 


aj   d 


_Xo< 

S22c 


qj 


oooooooooooo 
006006060066 

crjcOfO^^^M     N    >O«N    <N    C*5 


OOOOOOOOOOOO    ONOt^^  O>00 


oOO 


1  +  1  ++  1  +  1  +  1    1  + 


OOOOOOOOOOOO 


O   O    O   v>  ^o  O 
OOOOOOOOOOOO 


4-     +  1  ++  1       1  ++ 


O^O  O^iow  loiorot^-N  toO 
O  w  OOO\O  IOVOM  IOIOM 
l^t^f^O  O  O  rOfO^fOfOt^ 
OOOMMMOOwOOO 

666666666666 


OOOWMMOOMOOO 
666666666666 


OOOOOOOOOOOO 

+  1  ++      1    1  +  1    1    1    1 


MHHMMMVOUDMlOCOfOI-l 

666666666666 


00  00  00  00  00    O  O\00 

MMMMMIOIOMIOCOPOM 
MMMMMOOMO^S^M 

666666666666 


IRON  513 

brown  ferric  hydroxide,  is  in  large  measure  obviated,  the  blue  of 
the  starch  iodide  being  quite  clear  against  the  pale  green  chromic 
hydroxide.  After  standing  for  from  fifteen  to  twenty  minutes  the 
solution  is  made  alkaline  with  sodium  bicarbonate  and  an  excess 
of  standard  iodine  solution  added.  This  is  allowed  to  stand  in  a 
stoppered  flask  for  from  one-half  hour  to  an  hour,  the  excess  of 
iodine  being  then  removed  with  arsenious  acid  and  the  solution 
titrated  to  color  with  iodine  after  the  addition  of  starch. 

According  to  Edgar's  procedure  the  first  step  in  the  calculation 
is  the  reduction  to  terms  of  tenth  normal  solution  of  the  figures 
of  titration  obtained  in  the  processes  (I),  (II)  and  (III).  It  is 
evident  that  the  subtraction  of  the  titration  figure  of  (III)  from 
that  of  (I)  gives  the  number  of  cubic  centimeters  corresponding 
to  reduction  of  the  vanadium  pentoxide  to  tetroxide,  while  the 
subtraction  of  this  result  from  that  of  (II)  gives  the  number 
equivalent  to  the  reduction  of  the  ferric  salt.  By  multiplying 
these  figures  by  the  amounts  of  vanadic  acid,  chromic  acid  and 
ferric  oxide  equivalent  to  I  cm.3  of  the  n/io  reagent  the  weights 
of  these  substances  present  are  obtained.  An  example  of  this 
procedure  is  given  below: 

Titratiom.          cm.8  H/IO  factor.        cm.* 

(I)  31.15  Xi.ioo   =34.26  (tt/io)^V2O5+CrO,. 

(II)  23.4  Xi.ioo  =25.74  (w/io)~V2O5+Fe2O,. 

(Ill)  30.00—  8.74X1.000  =21.26  (n/io)~CrO3. 

34.26  —  21.26  =13.00  =3=V2O5. 

25.74—13.00  =12.74  =  Fe2O3. 

21 . 26X0.003334  (factor  for  CrO3)  =0.0709  grm.  CrO3  found. 
13.00X0.00912  (factor  for  V2O6)  =0.1185  grm.  V2O6  found. 
12.74X0.00799  (factor  for  Fe2O3)  =  0.1018  grm.  Fe2O3  found. 

Results  obtained  by  this  process  are  given  in  the  table  on  the 
preceding  page.  The  calculation  given  above  is  that  of  the  first 
determination  of  the  table. 


INDEX   OF   AUTHORS* 

PAGE 

Ashley,  R.  Harmon.     Dithionic  acid  and  dithionates,  determination  of.  369 

Sulphites,  estimation  of,  in  alkaline  solution 366 

Austin,  Martha.     Arsenic,  determination  of,  as  magnesium  pyroarsenate  288 

Beryllium,  not  determinable  as  pyrophosphate 153 

Cadmium,  estimation  of,  as  pyrophosphate 190 

Magnesium,    determination    of,    as    pyrophosphate 

(with  F.  A.  Gooch) 156 

Manganese,  determination  of 

separated   as  carbonate 481 

as  oxide  (with  F.  A.  Gooch) 478 

by  chlorate  process  (with  F.  A.  Gooch)  487 

as  pyrophosphate  (with  F.  A.  Gooch) .  482 

as  sulphate  (with  F.  A.  Gooch) 477 

Phosphoric   acid,    determination   of,    as    magnesium 

pyrophosphate  (with  F.  A.  Gooch) 282 

Zinc,  estimation  of,  as  pyrophosphate 185 

Beyer,  F.  B.     Filtering  crucible   in  electrolytic  analysis    (with   F.   A. 

Gooch^ 13 

Lead,  electrolytic  determination  of  (with  F.  A.  Gooch) .  .  252 
Manganese,   electrolytic   determination  of    (with  F.   A. 

Gooch) 485 

Blake,  J.  C.     Bromates,  iodometric  estimation  of  (with  F.  A.  Gooch). .  471 

Gold,  red  colloidal 150 

Blumenthal,  Philip  L.     Barium  and  strontium,  detection  of,  associated 

with  calcium  and  lead  (with  P.  E.  Browning)  160 
Lead,  detection  of,  in  sulphates  (with  P.  E. 

Browning) 252 

Bosworth,  Rowland  S.     Silver,  gravimetric  determination  of,  as  chro- 

mate  (with  F.  A.  Gooch) 136 

iodometric  determination  of,  precipitated 

as  chromate  (with  F.  A.  Gooch). . . .  140 
iodometric  determination  of,  reduced  by 

arsenite 143 

Boynton,  C.  N.     Barium,  estimation  of,  precipitated  by  acetyl  chloride 

in  acetone  (with  F.  A.  Gooch) 175 

separation  of,  from  calcium  and  magnesium 

(with  F.  A.  Gooch) 177 

Breckenridge,  J.  E.     Perchloric  acid,  preparation  of  (with  D.  A.  Kreider)  76 

Sodium,  detection  of  (with  D.  A.  Kreider) 74 

Brooks,  F.  T.     Chlorine,  bromine  and  iodine,  detection  of  (with  F.  A. 

Gooch) 440 

Brown,  James.     Oxidations  by  permanganate  in  presence  of  chlorides  .  .  52 
Browning,  Philip  E.     Arsenic  acid,  iodometric  estimation  of  (with  F.  A. 

Gooch) 291 

Arsenic,    antimony    and  tin,  estimation    of,   by 
ferricyanide    and   permanganate    (with    H.    E. 

Palmer) 322 

Barium,  precipitation  of,  as  sulphate 170 

Barium  and  calcium,  estimation  of,  separation  by 

amyl  alcohol  on  nitrates 166 

*  Workers  in  the  Kent  Chemical  Laboratory. 
515 


516  INDEX  OF  AUTHORS 

PAGE 

Browning,  Philip  E.     Barium  and  strontium, detection  of,  associated  with 

calcium   and   lead    (with 

P.  L.  Blumenthal) 160 

separation    of,     by    amyl 

alcohol  on  bromides.  .  .     167 
Barium,   strontium  and  calcium,   separation   of, 

by  amyl  alcohol  on  nitrates 162 

Barium  with  strontium,  and  calcium,  separation 

of,  and  estimation  by  amyl  alcohol  on  nitrates. .      166 
.         Cadmium,  estimation  of,  as  oxide    (with   L.    C. 

Jones) 188 

Caesium    and    rubidium,   estimation   of,   as   acid 

sulphates 106 

Cerium,  estimation  of,  by  ferricyanide  and  per- 
manganate (with  H.  E.  Palmer) 249 

separation  of,  from  cerium  earths  (with 

E.  J.  Roberts) 244 

iodometric  estimation  of,  digestion  process 
'(with  Hanford 
and  Hall)..  .  .     246 
distillation  proc- 
ess      (with 
Hanford   and 

Hall) 247 

Cerium  oxalate,  estimation  of,  by  permanganate 

(with  L.  A.  Lynch) 248 

Chromic  acid,  iodometric  determination  of 407 

Copper,  determination  of,  as  cuprous  iodide 114 

separation    from    cadmium    as    cuprous 

iodide 115 

Ferricyanides,  detection  of  (with  H.  E.  Palmer) . .     275 
Ferrocyanides,  detection  of  (with  H.  E.  Palmer) .     275 

Fluorine,  detection  of 432 

Iodine,    iodometric   determination   of,   in   haloid 

salts  (with  F.  A.  Gooch) 457 

Lead,  detection  of,  in  sulphates  (with  P.  L.  Blu- 
menthal)       252 

Magnesium,   separation   of,   from   alkalies   (with 

W.  A.  Drushel) 158 

Nickel,  detection  of,  in  presence  of  cobalt  (with 

J.  B.  Hartwell) . 491 

Potassium,  estimation  of,  as  pyrosulphate 92 

Silicon,  detection  of,  in  silicates  and  fluosilicates     241 

Sodium,  estimation  of,  as  pyrosulphate 79 

Strontium  and  calcium,  detection  of 163 

separation  of,  and  esti- 
mation       166 

Sulphides,  sulphates,  sulphites  and  thiosulphates, 

detection  of  (with  Ernest  Howe) 363 

Sulphocyanates,  detection  of  (with  H.  E.  Palmer)     276 
Tellurium,  separation  of,  from  selenium  (with  W. 

R.  Flint) 402 

Tellurium  dioxide,  precipitation  of  (with  W.  R. 

Flint) 402 

Thallium,  determination  of,  as  sulphates 219 

estimation  of,  by  precipitation  as  thallic 
hydroxide  (with  H.  E. 

Palmer) 220 

by  ferricyanide  and  per- 
manganate (with  H.  E. 
Palmer) 223 


INDEX  OF  AUTHORS  517 

PAGE 
Browning,  Philip  E.    Thallium,  determination  of,  as  chromate  (with  G. 

P.  Hutchins) 221 

iodometrically    (with    G. 

P.  Hutchins) 222 

Vanadium,  estimation  of.  as  silver  vanadate  (with 

H.  E.  Palmer) 328 

Vanadic  acid,  ipdometric  estimation  of,  reduced 

by  organic  acids  (with  R.  J.  Goodman) 341 

by  hydriodic  acid 343 

by  hyrlrobromic  acid 345 

Volatile  products,  removal  of,  without  loss  of  non- 
volatile material  (with  F.  A.  Gooch) 6 

demons,  C.  F.     Selenious   acid,    determination  of,   by   permanganate 

(with  F.  A.  Gooch) 382 

Curtis,  F.  W.     Vanadic  acid,  reduction  of,  by  hydrochloric  acid  (with  F.  A. 

Gooch) 334 

hydrobromic  acid  (with  F.  A. 

Gooch) 335 

hydriodic    acid    (with   F.   A. 

Gooch) 337 

Danner,  E.  W.     Antimony,  separation  of,  from  arsenic,  and  estimation 

(with  F.  A.  Gooch) .  „ . 311 

Oxygen,  loss  of,  in  oxidations  by  permanganate  (with 

F.  A.  Gooch) 43 

Tellurous  Acid,   determination  of,   by   permanganate 

(with  F.  A.  Gooch) 394 

Drushel,  W.  A.     Lanthanum,  estimation  of,  precipitated  as  oxalate.  .  .  .     218 
Magnesium,  separation  of,  from  alkalies  (with  P.  E. 

Browning) 158 

Potassium,  estimation  of,  as  cobalti-nitrite 93 

in  the  pure  salt 94 

in  mixtures  of  salts. ...       95 

in  fertilizers 96 

in  soils 97 

in  animal  fluids 98 

Eddy,  Ernest  A.     Magnesium,  determination  of,  as  oxide  (with  F.  A. 

Gooch) . . 154 

Edgar,  Graham.     Chromic  acid  and  vanadic  acid,  estimation  of 409 

Iron  and  vanadium,  estimation  of 508 

Ferric  iron,  vanadic  acid  and  chromic  acid,  estima- 
tion of 510 

Molybdic  acid  and  vanadic  acid,  determination  of,  by 

reductions  and  oxidations 427 

Vanadic  acid  and  antimonic  acid,  estimation  of.  ...     350 

and  arsenic  acid,  estimation  of 350 

.     reduction  of,  by  zinc,  with  use  of  ferric 

sulphate  (with  F.  A.  Gooch) 349 

Volatile  products,  distillation  and  absorption  of 5 

Ensign,  J.  R.     Bromine  (and  chlorine),  determination  of,  in  mixtures  of 
alkali  bromides  (and  chlorides)  with  iodides  (with  F.  A. 

Gooch) 452 

Evans,  P.  S.  Jr.     Selenic  acid,  iodometric  determination  of,  reduced  by 

hydrochloric  acid  (with  F.  A.  Gooch) 385 

Fairbanks,  Charlotte.     Halogens,  determination  of,  by  electrolytic  re- 
duction of  silver  salts  (with  F.  A.  Gooch) ....     459 
Molybdic  acid,  iodometric  estimation  of  (with 

F.  A.  Gooch) 415 

distillation  process  (with  F. 

A.  Gooch) 416 

reoxidation  of  residue  (with 
F.  A.  Gooch) 420 


518  INDEX  OF  AUTHORS 

PAGR 

Fairbanks,  Charlotte.     Phosphorus,  iodometric  determination  of,  in  iron     283 
Feiser,  J.  P.     Silver,  electrolytic  determination  of,  in  oxalate  solution 

(with  F.  A.  Gooch) 138 

Flint,  William  R.     Tellurium,  separation  of,  from  selenium  (with  P.  E. 

Browning) 402 

Tellurium    dioxide,    precipitation    of    (with    P.    E. 

Browning) 402 

Flora,  Charles  P.     Cadmium,  estimation  of,  as  oxide 188 

electrolytic  estimation  of 191 

from  sulphuric  acid  solution 191 

from  solutions  of  acetates 192 

from  solutions  of  cyanides 193 

from  solutions  of  pyrophosphates  and 

orthophosphate 194 

Gilbert,  R.  D.     Ammonium  vanadate,  precipitation  of,  by  ammonium 

chloride  (with  F.  A.  Gooch) 326 

Vanadic  acid,  estimation  of,  by  Jones  reductor  and 

silver  sulphate  (with  F.  A.  Gooch) 346 

Gillespie,  David  H.  M.     Alkali  hydroxide,  determination  of,  by  reaction 

with  iodine  (with  C.  F.  Walker) 71 

Gooch,  F.  A.     Aluminium,  estimation  of,  by  precipitation  as  chloride 

(with  F.  S.  Havens) 214 

separation  of,  from  iron  (with  F.  S.  Havens)     214 
Aluminium  sulphate,  hydrolysis  of,  in  bromide-bromate 

mixture  (with  R.  W.  Osborne) 70 

Antimonic  acid  and  arsenic  acid,  iodometric  determina- 
tion of  (with  H.  W.  Gruener) 308 

Antimony,  separation  of,  from  arsenic,  and  estimation 

(with  E.  W.  Danner) 311 

Arsenic,  antimony  and  tin  detection  of  (with  B.  Hodge) .     312 

(with  I.  K.  Phelps).     316 
Arsenic    acid,  iodometric  estimation  of    (with    P.    E. 

Browning) 291 

(with  J.  C.  Morris) 294 

estimation  of  small  amounts,  precipitated 
as     ammonium     magnesium     arsenate 

(with  M.  A.  Phelps)- 290 

estimation  of  minute  amounts,  in  copper 

(with  H.  P.  Moseley) 301 

separation  of,  from  copper,  as  ammonium 

magnesium  arsenate  (with  M.  A.  Phelps)     303 
Barium,  estimation  of,  as  chloride  (with  C.  N.  Boynton)     175 
separated  from  calcium  and  magnesium  (with 

C.  N.  Boynton) 177 

Boric  acid,  gravimetric  determination   of    (with   L.    C. 

Jones) 201 

with  sodium  tungstate  as  re- 
tainer (with  L.  C.  Jones) .  .     204 

Bromates,  iodometric  estimation  of  (with  J.  C.  Blake) .  .     471 
Bromine  (and  chlorine),  determination  of,  in  mixtures  of 
alkali  bromides  (and  chlorides)  with  iodides  (with  y.  R. 

Ensign) " 452 

Carbon  dioxide,  determination  of,  by  ignition  (with  S.  B. 

Kuzirian) 226 

precipitation  and  gravimetric  estimation        • 

of  (with  I.  K.  Phelps)    228 

Chlorates,  iodometric  estimation  of  (with  C.  G.  Smith) .  .     463 
Chlorine,  determination  of.  in  mixtures  of  alkali  chlorides 

and  iodides  (with  F.  W.  Mar) 449 

fixation  of,  on  silver  anode  (with  H.  L.  Read) .  .       20 


INDEX  OF  AUTHORS  519 

PAGE 
Gooch,  F.  A.     Chlorine,  bromine  and  iodine,  detection  of  (with  F.  T. 

Brooks) 440 

Chromium,  estimation  of,  as  silver  chromate  (with  L.  H. 

Weed) 406 

Copper,  determination  of,  by  titration  of  oxalate  (with 

H.  L.  Ward) .....".. 126 

electrolytic  determination  of,  by  rotating  cathode 

(with  H.  E.  Medway) . . 116 

iodometric  estimation  of  (with  F.  H.  Heath) ....      121 
Filtering  crucible   in  electrolytic  analysis   (with  F.   E. 

Beyer) 13 

Gaseous  products,  evolution  of,  without  mechanical  loss 

(with  C.  F.  Walker) 6 

Halogens,  determination  of,  by  electrolytic  reduction  of 

silver  salts  (with  C.  Fairbanks) 459 

Iodides,  analysis  of,  by  iodic  acid  (with  C.  F.  Walker) .  .  .     454 
gravimetric  determination  of,  by  absorption  in 

silver  (with  C.  C.  Perkins) 444 

Iodine,    iodometric    determination    of,    in    haloid    salts 

(with  P.  E.  Browning) 457 

standardization  of,  by  silver  (with  C.  C.  Perkins)       27 
Iron,  estimation  of,  in  presence  of  titanium  (with  H.  D. 

Newton) .     500 

separations  of,  by  volatilization  in  hydrogen  chlo- 
ride (with  F.  E.  Havens) 504 

Lead,  electrolytic  determination  of  (with  F.  B.  Beyer). .     252 
Magnesium,  determination  of,  as  oxide  (with  E.  A.  Eddy)     154 

as     pyrophosphate     (with 

M.  Austin) 156 

Manganese,  determination  of,  by  chlorate  process  (with 

'  M.  Austin) 487 

by  electrolysis  (with  F.  B. 

Beyer)... .     485 

as  oxide  (with  M.Austin)     478 
as    pyrophosphate    (with 

M.  Austin) 482 

as  sulphate  (with  M.  Aus- 
tin)      477 

Molybdic  acid,  iodometric  estimation  of, 

by  distillation  process  (with  C.  Fairbanks) 416 

(with  J.  T.  Norton,  Jr.) 418 

by  iodine  oxidation  of  residue  (with  C.  Fairbanks) ....     416 
by  permanganate  oxidation  of   residue,   (with   O.    S. 

Pulman,  Jr.) 422 

Nitrates,  estimation  of,  by  ignition  (with  S.  B.  Kuzirian)     256 
iodometric  estimation  of  (with  H.  W.  Gruener), 

263,  266,  268 
Nitrates    and    chlorates,    estimation    of    (with    H.    W. 

Gruener) 273 

Oxygen,  loss  of,  in  permanganate  titrations  (with  L.  W. 

Danner) 43 

(with  C.  A.  Peters)    . .48,  50 

Perchlorates,  detection  of,  in  association  with  chlorides, 

chlorates  and  nitrates  (with  D.  A.  Kreider) 465 

Phosphoric  acid,  determination  of,  as  magnesium  pyro- 
phosphate (with  M.  Austin) 282 

Potassium,  spectroscopic  detection  and  determination  of 

(with  T.  S.  Hart) 80 

Precipitates,  purification  of 10 

Rotating  cathode  (with  H.  E.  Medway) 1 1 


520  INDEX  OF  AUTHORS 

PACK 
Gooch,  F.  A.     Rubidium,  spectroscopic  determination  of   (with  J.   I. 

Phinney) 102 

Selenic  acid,  iodometric  determination  of, 

reduction  by  hydrochloric  acid  (with  P.  S.  Evans,  Jr.)     385 

by  hydrobromic  acid  (with  W.  S.  Scoville)..     386 

by  hydriodic  acid  (with  VV.  C.  Reynolds)..  .     388 

by  differential  method  (with  A.  W.  Peirce). .     380 

Selenious  acid,  determination  of,  by  permanganate  (with 

C.  F.  demons) 382 

iodometric  determination  of  (with  W.  G. 

Reynolds). 378 

by   differential  method 

(with  A.  W.  Peirce) .     380 
Selenium  and  tellurium,  separation  of,  by  difference  in 

volatility  of  bromides  (with  A.  W.  Peirce) 390 

Silver,  gravimetric  determination  of,  as  chromate  (with 

R.  S.  Bosworth) 136 

electrolytic  determination  of,  in  oxalate  solution 

(with  J.  P.  Feiser) 138 

iodometric  determination  of,  as  chromate  (with  R. 

S.  Bosworth) 140 

Telluric  acid,  iodometric  determination  (with  J.  How- 
land)  401 

Teliurous  acid,  determination  of,  by  permanganate  (with 

E.  W.  Danner)     394 
in  presence  of  bro- 
mide   (with   C. 

A.  Peters) 397 

in      presence      of 
chloride      (with 
C.A.Peters)..     396 
by     precipitation      as 
iodide  (with  W. 
C.  Morgan).  . .     398 
iodometric  (with  C.  A. 

Peters) 399 

Vanadic  acid,  estimation  of,  precipitated  by  ammonium 

chloride  (with  R.  D.  Gilbert).  ........     326 

reduced  by  hydrochloric  acid 

(with  L.  B.  Stookey) 330 

by  hydrochloric  acid 
(with  R.  W.  Curtis) .     334 
by  hydrobromic  acid 

(with  R.W.  Curtis)     335 
by     hydriodic     acid 

(with  R.W.  Curtis)     337 
by  Jones  reductor,  with  silver  sul- 
phate (with  R.  D.  Gilbert) 346 

by  zinc,  with  ferric  sulphate  (with 

Graham  Edgar) 349 

Volatile  products,  removal  of,  without  loss  of  non- 
volatile material  (with  P.  E.  Brown- 
ing)    <S 

distillation  of,  and  absorption   (with 

J.  T.  Norton,  Jr.) 4 

(with  A.  W.  Peirce) 5 

Goodman,  Richard  J.     Vanadic  acid,  iodometric  estimation  of,  reduced 

by  organic  acids  (with  P.  E.  Browning) 341 

Gruener,  H.  W.     Antimonic  acid  and   arsenic  acid,  iodometric  deter- 
mination of  (with  F.  A.  Gooch) 308 


INDEX  OF  AUTHORS  521 

PAGE 
Gruener,  H.  W.     Nitrates,    iodometric    estimation     of     (with    F.    A. 

Gooch) 263,  266,  268 

Nitrates   and   chlorates,   estimation  of    (with  F.   A. 

Gooch) 273 

Standard  tartar  emetic 38 

Hale,  F.  E.     Standard  tartar  emetic 39 

Starch  indicator  for  free  iodine 29 

Hall,  F.  J.      Cerium,  iodometric  estimation  of  (with  P.  E.  Browning) 

246,  247 

Hanford,  G.  A.     Cerium,  iodometric  estimation  of  (with  P.  E.  Brown- 
ing)   -246,  247 

Hart,  T.  S.     Potassium,  spectroscopic  detection  and  determination  of 

(with  F.  A.  Gooch) 80 

Hartwell,  John  B.     Nickel,  detection  of,  in  presence  of  cobalt  (with  P.  E. 

Browning) 491 

Havens,  Franke  Stuart.     Aluminium,   determination  of,   by   precipita- 
tion    with    ether-hydrochloric 

acid  (with  F.  A.  Gooch) 214 

separation  of,  from  bismuth  and 

mercury 217 

and  beryllium,  determination  of . .     216 
and  copper,  determination  of.  ...     217 

and  zinc,  determination  of ......     216 

and  iron,  separation  of  (with  F.  A. 

Gooch) 214 

Beryllium  chloride,  conversion  of,  to  oxide ....      153 
Beryllium   oxide,    separation   of,    from   ferric 

oxide  (with  A.  F.  Way) 154 

Nickel  and  cobalt,  separation  of,  by  ethereal 

hydrochloric  acid 492 

Iron,  separations  of,  by  volatilization  in  hydro- 
gen chloride 504 

from  aluminium  (with  F. 

A.  Gooch) ; . .  .     506 

from   beryllium    (with  A. 

F.Way) 507 

from  chromium   (with  A. 

F.  Way) 507 

from   zirconium    (with  A. 

F.Way) 508 

Heath,  F.  H.     Arsenic,  antimony  and  copper,  determination  of 318 

Copper,  iodometric  determination  of  (with  F.  A.  Gooch)     121 
Hileman,  Albert.     Fluorine,  estimation  of,  evolved  as  silicon  fluoride .  .  .     436 

iodometric  determination  of,  in  fluorides .     439 

Fluosilicic  acid,  acidimetric  estimation  of 432 

iodometric  estimation  of 435 

Hodge,  B.     Arsenic,  antimony  and  tin,  detection  of  (with  F.  A.  Gooch)     313 
Howe,  Ernest.     Sulphides,  sulphates,  sulphites  and   thiosulphates,  de- 
tection of  (with  P.  E.  Browning) 363 

Howland,  J.     Telluric  acid,  iodometric  determination  of   (with  F.  A. 

Gooch) 401 

Hubbard,  J.  L.     Succinic  acid,  use  of,  as  a  standard  in  neutralization 

processes  (with  I.  K.  Phelps) 54 

Hutchins,  George  P.     Thallium,  gravimetric  estimation  of,  as  chromate 

(with  P.  E.  Browning) 221 

iodometric  estimation  of  (with  P.  E. 

Browning) 222 

Jones,  L.  C.     Boric  acid,  acidimetric  estimation  of 205 

gravimetric    determination     of    (with    F.    A. 

Gooch) . .  .     201 


522 
Jones,  L.  C. 


INDEX  OF  AUTHORS 


PAGE 


88 
74 

7 

225 
i 

2 


Boric  acid,  gravimetric  determination  of,  with    sodium 

tungstate 
as  a  retainer 
(with  F.  A. 
Gooch). . .  .  204 

iodometric  determination  of 210 

neutralization  of  acids  associated  with 206 

strengthening  of,  by  mannite 208 

Cadmium,  estimation  of,  as  oxide  (with  P.  E.  Browning) .      188 

Kreider,  D.  Albert.     Force  pump & 

Oxygen,  iodometric  estimation  of,  in  air 355 

water 355,  360 

Perchlorates,    detection    of,    in    association    with 
chlorides,   chlorates    and    nitrate 

t  (with  F.  A.  Gooch) 465 

iodometric  determination  of 467 

Perchloric  acid,  preparation  of 88 

(with  J.  E.  Breckenridge) 76 

Potassium,   separation  of,  and  determination  as 

perchlorate  

Sodium,  detection  of  (with  J.  E.  Breckenridge) .... 

Valve 

Kreider,  J.  Lehn.     Carbon  dioxide,  determination  of,  in  carbonates,  by 

loss 

Gaseous  products,  determination  of,  by  loss 

Hydrogen,  determination  of,  by  loss 

Nitrogen,  determination  of,  in  ammonia  compounds 

and  derivatives 256 

Kuzirian,  S.  B.     Carbon  dioxide,  determination  of,  by  ignition  (with  F. 

A.  Gooch) 226 

Nitrates,  estimation  of,  by  ignition  (with  F.  A.  Gooch)     256 
Lynch.  Leo  A.     Cerium  oxalate,  estimation  of,  by  permanganate  (with 

P.  E.  Browning) 248 

Mar,  F.  W.     Barium,  estimation  of,  as  chloride 174 

as  sulphate,  in  presence  of  hydro- 
chloric acid 168 

separation  of,  as  chloride,  from  calcium  and  mag- 
nesium        175 

Barium  sulphate,  purification  of 172 

Chlorine,  determination  of,  in  mixtures  of  alkali  chlorides 

and  iodides  (with  F.  A.  Gooch) 449 

Maryott,  C.  H.     Halogens,  determination  of,  in  benzol  derivatives,  by 

the  use  of  metallic  potassium 447 

Maxson,  Ralph  Nelson.     Gold,  colorimetric  determination  of 150 

iodometric  determination  of 150 

Medway,  H.  E.     Cadmium,  electrolytic  determination  of 191 

Copper,   electrolytic    determination   of    (with    F.    A. 

Gooch) 116 

Gold,  electrolytic  determination  of 145 

Silver,  electrolytic  determination  of 117 

Tin,  electrolytic  determination  of 251 

Zinc,  electrolytic  determination  of 186 

Rotating  cathode  (with  F.  A.  Gooch) 1 1 

Moody,  Seth  E.     Hydrolysis  of  salts  in  presence  of  iodide-iodate  mix- 
tures         6 1 

Persulphates,  determination  of  (with  C.  A.  Peters) ...     370 
Morgan,  W.  C.     Tellurous  acid,  determination  of,  by  precipitation,  as 

tellurous  iodide  (with  F.  A.  Gooch) 398 

Morley,  Frederick  H.     Gold,  iodometric  estimation  of,  small  amounts 

(with  F.  A.  Gooch) 146 


INDEX  OF  AUTHORS  523 

PAGE 
Morris,  Julia  C.     Arsenic  acid,  iodometric  estimation  of   (with  F.  A. 

Gooch) 294 

Moseley,  H.  P.     Arsenic,   estimation  of,   minute  quantities  in  copper 

(with  F.  A.  Gooch) 301 

Newton,  H.  D.     Iron,  estimation  of,  in  presence  of  titanium  (with  F.  A. 

Gooch) 500 

by  permanganate,  after  reduction 

with  titanous  sulphate 503 

Titanic  acid,  determination  of,  by  reduction  and  per- 
manganate titration 242 

Norton,  John  T.  Jr.     Iron,  determination  of,  in  ferric  state,  by  thiosul- 

phate  and  iodine , 492 

Mercury,    determination    of,    by    titration    with 

thiosulphate 196 

Molybdic  acid,  iodometric  estimation  of  (with  F. 

A.  Gooch) 418 

Selenious  acid,  determination  of,  by  the  method  of 

Norn's  and  Fay 383 

Thiosulphates,  iodometric  determination  of  (with 

F.  A.  Gooch) -.-•.•••: 364- 

Volatile  products,  distillation  and  absorption  of 

(with  F.  A.  Gooch)     4 

Osborne,  R.  W.     Aluminium   sulphate,   hydrolysis   of,    in   presence   of 

bromide-bromate  mixture  (with  F.  A.  Gooch) ....        70 
Palmer,  H.  E.     Arsenic,   antimony  and  tin,  estimation  of,  by    ferri- 
cyanide and  permanganate  (with  P.  E.  Browning) .  .     322 
Cerium,  estimation  of,  by  ferricyanide  and  perman- 
ganate (with  P.  E.  Browning) 249 

Chromium  in  chromic  condition,  estimation  of 413 

Chromic  acid  and  yanadic  acid,  estimation  of,  by  reduc- 
tions and  oxidations 411 

Ferricyanides,  detection  of  (with  P.  E.  Browning)  ....  275 
Ferrocyanides,  detection  of  (with  P.  E.  Browning)  ....  275 
Sulphocyanates,  detection  of  (with  P.  E.  Browning)  ....  276 
Thallium,  estimation  of,  as  thallic  hydroxide  (with  P.  E. 

Browning) 220 

by    ferricyanide    and    perman- 
ganate (with  P.  E.  Browning)     223 
Vanadium,  estimation  of,  as  silver  vanadate  (with  P.  E. 

Browning) 328 

Vanadium  tetroxide,  estimation  of,  by  ferricyanide  and 

permanganate 352 

Peirce,  A.  W.     Selenic  acid,  iodometric  determination  of,  by  differential 

method  (with  F.  A.  Gooch)      388 

Selenious  acid,  gravimetric    determination    of,    by    pre- 
cipitation of  selenium 376 

iodometric  determination  of,  by  differen- 
tial method  (with  F.  A.  Gooch) 380 

Selenium  and  tellurium,  separation  of,  by  difference  in 

volatility  of  the  bromides  (with  F.  A.  Gooch) 390 

Volatile  products,  distillation  of,  and  absorption  (with 

F.  A.  Gooch) 5 

Perkins,  Claude  C.     Determination   of   iodine  by  absorption   in   silver 

(with  F.  A.  Gooch) 444 

Determination  of  oxidizers  by  liberation  of  iodine  and 

absorption  of  that  element  by  silver 361 

chlorine  and  bromine 443 

molybdic  acid 415 

selenious  acid 375 

tellurious  acid 394 


.324 

Perkins,  Claude  C. 


INDEX  OF  AUTHORS 


PAGE 

Determination  of  vanadic  acid 325 

Rotary  shaker 9 

Standardization  of  iodine  by  silver(with  F.  A.  Gooch)       27 

Standardization  of  permanganate 42 

Peters,  Charles  A.     Calcium,  strontium  and  barium,  determination  of, 

as  oxalates,  by  permanganate 181 

Copper,  determination  of,  by  permanganate  titra- 

tion  of  oxalate 123 

separation  of,  as  oxalate,  from  cadmium 

arsenic,  iron,  tin  and  zinc 131 

Electrolysis  of  sodium  chloride  with  silver  anode  and 

mercury  cathode ;•.-••. 22 

Mercury,  estimation  of,  by  precipitation  by  ammo- 
nium oxalate  and  titration  of  the  excess     197 
gravimetric  determination  of,  as  oxalate     195 
Oxygen,  loss  of,  in  permanganate  titrations  in  pres- 
ence of  hydrochloric  acid  (with  F.  A.  Gooch). ...          50 
Persulphates,  determination  of  (with  S.  E.  Moody)     370 
Strontium  and  barium,  gravimetric  determination 

of,  as  oxalates 181 

Tellurous  acid,  determination  of,  by  permanganate, 
in    presence     of     bromide     (with 

F.  A.  Gooch) 397 

in  presence  of  chloride  (with  F.  A. 

Gooch) 396 

iodometric  estimation   (with  F.  A. 

Gooch) 399 

Titrations  by  permanganate  in  presence  of  hydro- 
chloric acid  (with  F.  A.  Gooch) 49 

Phelps,  I.  K.     Alkali  hydroxides,  determination  of,  by   reaction  with 

iodine 71 

Alkali  hydroxides  and  carbonates,  iodometric  determina- 
tion of  (with  L.  H.  Weed) 60 

Arsenic,  antimony  and   tin,    detection   of    (with  F.  A. 

Gooch) 316 

Carbon,  determination  of,  in  organic  substances, 

by  oxidation  with  chromic  acid 236 

by  oxidation  with  permanganate '. 234 

Carbon  dioxide,  iodometric  determination  of .  .  .  , 231 

in  carbonates.  •  v  •  • ; .-  •' 232 

gravimetric  determination  of  (with  F.  A.  "P 

Gooch) 228 

Chlorates,  estimation  of,  by  reduction  with  ferrous  sul- 
phate      462 

Nitrates,  estimation  of,  by  reduction  with  a  ferrous  salt 

and  titration 258 

Nitrites,  iodometric  estimation  of 269 

Organic  acids  and  acid  anhydrides,  use  of,  as  standards, 

in  neutralization  processes  (with  L.  H.  Weed) 56 

in  iodometric  processes  (with  L.  H.  Weed) 56,  59 

Organic  substances,  combustion  of,  in  the  wet  way 234 

Oxygen,  determination  of,  in  organic  substances 239 

Succinic  acid,  use  of,  as  a  standard  in  neutralization 

processes  (with  J.  L.  Hubbard) 54 

Phelps,  M.  A.     Arsenic  acid,  estimation  of,  small  amounts,  precipitated 

as  ammonium  magnesium  arsenate  (with 

F.  A.  Gooch) 290 

separation  of,  from  copper,  as  am- 
monium magnesium  arsenate  (with 
F.  A.  Gooch) 3<>5 


INDEX  OF  AUTHORS  525 

PAGE 

Phinney,  J.  I.     Barium  sulphate,  purification  of 172 

Rubidium,  spectroscopic  determination  of  (with  F.  A. 

Gooch) 102 

Pulman,  O.  S.,  Jr.  Molybdic  acid,  iodometric  estimation  of,  by  reduc- 
tion with  hydriodic  acid  and  reoxidation  by  per- 
manganate (with  F.  A.  Gooch) ^ 422 

Phosphoric    acid,    estimation    of,    precipitated    as 

uranyl  phosphate 286 

Uranium,   determination   of,   by  aid   of  the  Jones 

reductor 430 

Randall,  D.  L.     Ferric  chloride,  behavior  of,  in  Jones  reductor 497 

Ferrous  sulphate,  permanganate  titration  of,  in  presence 

of  nitric  acid 499^ 

Mercurous  salts,  titration  of,  by  permanganate 198 

Molybdic  acid,  reduction  of,  in  Jones  reductor. 424 

Phosphoric  acid,  volumetric  estimation  of,  precipitated 

as  ammonium  phosphomolybdate 285 

Read,  H.  L.     Fixation  of  chlorine  on  silver  anode  (with  F.  A.  Gooch) .  .  20 
Reynolds,  W.  G.     Selenic  acid,  iodometric  determination  of  (with  F.  A. 

Gooch) 388 

Selenious  acid,  iodometric  determination  of  (with  F. 

A.  Gooch) 378 

Roberts,  Charlotte  F.     Nitrates,  estimation  of 260 

Nitrates  and  chlorates,  estimation  of 273. 

Nitrates  and  nitrites,  estimation  of 271 

Nitrites,  estimation  of 271 

Permanganate  standardization,  in  iron  analysis  495 
Roberts,  Edwin  J.     Cerium,  separation  of,  from  cerium  earths  (with  P. 

E.  Browning) 244, 

Scoville,  W.  S.     Selenic  acid,  iodometric  determination  of  (with  F.  A. 

Gooch) 386 

Smith,  C.  G.     Chlorates,  iodometric  determination  of  (with  F.  A.  Gooch)  463 
Stookey,  L.  B.     Vanadic  acid,  reduction  of,  by  hydrochloric  acid,  and 

estimation  (with  F.  A.  Gooch) 330 

Thorne,  Norman  C.     Barium,  separation  of,  as  bromide,  from  calcium 

and  magnesium 180 

Barium  bromide,  precipitation  of,  by  ether-hydro- 

bromic  acid . 1 79 

Van  Name,  R.  G.     Copper,  gravimetric  determination  of,  as  sulpho- 

cyanate. 108 

separation  of,  as  sulphocyanate,  from  bis- 
muth, antimony,  tin  and  arsenic 112 

Sulphocyanates,  gravimetric  determination  of 276 

volumetric  determination  of 279* 

Walker,  Claude  F.     Alkali  hydroxides,  determination  of,  by  reaction 

with  iodine  (with  D.  H.  M.  Gillespie) 70 

Gaseous  products,  removal  of,  without  mechani- 
cal loss  (with  F.  A.  Gooch) 6 

Iodides,  analysis  of,  by  iodic  acid  (with  F.  A.  Gooch)  454 
Ward,  H.  L.     Copper,  determination  of,  by  titration  of  oxalate  (with 

F.  A.  Gooch) 126 

associated  with  lead 135 

separated  from  cadmium,  arsenic  and  iron 132 

Lead,  estimation  of,  by  titration  of  oxalate 254, 

Nickel,  estimation  of,  by  titration  of  oxalate 490 

Zinc,  estimation  of,  by  titration  of  oxalate 187 

Way,  Arthur  Fitch.     Separations  by  volatilization  in  hydrogen  chloride, 

iron  and  beryllium  (with  F.  S.  Havens) 507 

iron  and  chromium  (with  F.  S.  Havens) 507 

iron  and  zirconium  (with  F.  S.  Havens) 508 


526  INDEX  OF  AUTHORS 

PAGE 

Weed,  L.  H.     Alkali  hydroxides  and  carbonates,  iodometric  determina- 
tion of  (with  I.  K.  Phelps) 60 

Chromium,  estimation  of,  as  silver  chromate  (with  F.  A. 

Gooch) 406 

Organic  acids  and  acid  anhydrides, 

as  standards  in  neutralization  processes   (with   I.  K. 

Phelps) 56 

as   standards    in    iodometric    processes    (with     I.    K. 

Phelps) , 56,  59 


INDEX   OF   SUBJECTS 

PAGE 

Acidimetry 54 

Acids,  determination  of,  by  neutralization .  . 54,  56 

by  iodide-iodate  mixture 59,  72 

liberated  in  hydrolysis 61 

standards  in  neutralization  processes. 56,  57,  58 

Acid  anhydrides,  as  standards  in  neutralization  processes 54»  59 

succinic  anhydride 57 

phthalic  anhydride 58 

Alkalimetry ^ 54 

Alkali  carbonates,  iodometric  determination  of 60 

Alkali  hydroxides,  determination  of,  by  neutralization 54,  56 

by  succinic  acid 54 

by  organic  acids 56 

by  acid  anhydrides 56 

iodometric,  by  iodide-iodate  mix- 
ture         59 

by  iodine 70 

Aluminium,  determination  of,  precipitated  by  ether-hydrochloric  acid. .     214 

separation  of,  from  beryllium 216 

bismuth 217 

copper 217 

iron 214 

mercury ...     217 

zinc 216 

separation  of,  from  iron,  by  gaseous  hydrogen  chloride 506 

Aluminium  sulphate,  hydrolysis  of,  in  bromide-bromate  mixture 70 

in  iodide-iodate  mixture 62 

Alums,  analysis  of,  determination  of  basic  alumina,  or  free  acid 67 

Ammonium  sulphate,  hydrolysis  of,  in  iodide-iodate  mixture 65 

Antimonic  acid,  iodometric  determination  of 308 

and  arsenic  acid,  estimation  of 308 

and  vanadic  acid,  estimation  of 325,  350 

Antimony,  detection  of,  by  hydrochloric  acid  and  bromide 316 

by  hydrochloric  acid  and  iodide 313 

estimation  of,  by  ferricyanide  and  permanganate,  associated 

with  arsenic  and  tin .322,  323 

iodometric,  associated  with  copper 318 

separation  of,  from  arsenic,  and  determination  of,  by  hydro- 
chloric acid  and  iodide 311 

from  copper  precipitated  as  sulphocyanate ....     1 12 

Arsenic,  detection  of,  by  hydrochloric  acid  and  bromide 316 

by  hydrochloric  acid  and  iodide 313 

determination  of,  as  magnesium  pyroarsenate 288 

small  amounts  precipitated  by  freezing..     290 

minute  quantities  in  copper 301 

by    ferricyanide    and     permanganate,    asso- 
ciated with  antimony  and  tin 322,  324 

iodometric,  associated  with  copper 318 

separation  of,  from  antimony,  by  hydrochloric  acid  and  iodide.     311 

from  copper  precipitated  as  oxalate.i3i,  133,  134,  135 

as  sulphocyanate 112 

527 


528  INDEX  OF  SUBJECTS 

PAGE 

Arsenic  acid,  estimation  of,  iodometric 291 ,  294 

reduction  by  hydriodic  acid,  with  titration 

of  liberated  iodine 295 

and  antimonic  acid,  estimation  of 308 

and  vanadic  acid,  estimation  of 325,  350 

Arsenic  trioxide,  use  of,  as  an  iodometric  standard 29 

in  standardizing  permanganate,  without  iodine.  411 

with  iodine. ...  412 

Barium,  detection  of,  associated  with  calcium  and  lead 160 

estimation  of,  as  sulphate  in  presence  of  hydrochloric  acid 168 

nitric  acid  and  aqua 

regia 170 

gravimetric  estimation  of,  precipitated  as  oxalate  from  alcohol  180 

precipitation  of,  by  acetyl-chloride  in  acetone,  and  estimation. .  175 

by  ether-hydrobromic  acid,  and  estimation ....  179 

by  ether-hydrochloric  acid,  and  estimation 174 

separation  of,  from  calcium,  and  estimation  by  action  of  amyl 

alcohol  on  the  nitrates 162,  166 

from  calcium  and  magnesium,  and  estimation, 

by  acetylchloride  in  acetone 177 

by  ether  hydrobromic  acid 180 

by  ether  hydrochloric  acid 175 

from  strontium,  and  estimation,  by  action  of  amyl 

alcohol  in  the  bromides 167 

volumetric  estimation  of,  precipitated  as  oxalate  from  alcohol . .  184 
Barium  with  strontium  and  calcium,  estimation  of,  by  action  of  amyl 

alcohol  on  the  nitrates 166 

Barium,  strontium  and  calcium,  separation  of,  by  action  of  amyl  alcohol 

on  the  nitrates 162 

Barium  sulphate,  purification  of,  after  precipitation 172 

Beryllium,  conversion  of  chloride  to  oxide 1 53 

not  determinable  by  precipitation  as  ammonium  beryllium 

phosphate : 153 

separation  from  aluminium,  by  ether-hydrochloric  acid.  ...  216 

from  iron,  by  gaseous  hydrogen  chloride 154,  507 

Bismuth,  separation  of,  from  aluminium,  by,  ether-hydrochloric  acid, 

and  determination 217 

from  copper  precipitated  as  sulphocyanate .  ...  112 

Boric  acid,  acidimetric  estimation  of 205 

neutralization  of  stronger  acids 206 

strengthening  of  acidity  by  mannite..  208 

gravimetric  estimation,  with  calcium  oxide  as  retainer 201 

with  sodium  tungstate  as  retainer .  . .  204 

iodometric  determination 210 

Bromates,  iodometric  estimation  of,  reduction  by  arsenate-iodide  mixture  475 

by  arsenious  acid 474 

by  hydriodic  acid 471 

Bromide-bromate  mixture,  reaction  of,  with  aluminium  sulphate 70 

Bromine,  determination  of,  in  benzol  derivatives 447 

gravimetric  determination  of,  by  liberation  of  iodine  and  ab- 
sorption of  that  element  by  silver 443 

Bromine  and  chlorine,  determination  of,  in  alkali,  bromides,  chlorides, 

and  iodides 452 

Bromine,  chlorine  and  iodine,  detection  of 440 

Cadmium,  electrolytic  determination  of, 

with  the  rotating  cathode 191 

from  solutions  of  acetates 192 

cyanides 193 


INDEX  OF  SUBJECTS  529 

PAGE 

Cadmium,  electrolytic  determination  of,  in  sulphuric  acid 191 

in    pyrophosphates   and   ortho- 
phosphates 194 

estimation  of,  as  oxide,  precipitated  as  carbonate 188 

as  hydroxide 189 

as  pyrophosphate 190 

separation  of,  from  copper  precipitated  as  oxalate 131 ,  133, 134 

Caesium,  estimation  of,  as  acid  sulphate 106 

Calcium,  detection  of,  with  strontium .  . . 163 

separation  of,  from  barium,  and  estimation,  by  the  action  of 

amyl  alcohol  on  the  nitrates 162,  166 

from  strontium,  and  estimation,  by  the  action  of 

amyl  alcohol  on  the  nitrates 162,  164 

from  strontium  and  barium,  and  estimation,  by 

amyl  alcohol  on  the  nitrates 162,  166 

volumetric  estimation  of,  precipitated  as  oxalate.  ••;••-. 181 

Carbon,  determination  of,  in  organic  substances,  by  combustion  in  the 

wet  way 234 

by  permanganate 234 

by  chromic  acid 236 

Carbon  dioxide,  determination  of,  in  carbonates,  by  loss  in  action  of  acid  225 

by  ignition  with  sodium  paratungstate  226 

iodometric 231,  232 

in  oxidations  by  permanganate,  iodo- 
metric   234 

in  oxidations  by  chromic  acid,  iodo- 
metric    236 

precipitation  of,  and  gravimetric  determination 228 

Cerium,  iodometric  estimation  of,  by  digestion  process 246 

by  distillation  process 247 

separation   of,   from   cerium   earths,   by   bromine   and   alkali 

hydroxide 244 

volumetric  estimation  of,  by  ferricyanide  and  permanganate .  .  249 

precipitated  as  oxalate 248 

Chlorates,  estimation  of,  iodometric 463 

reduced  by  ferrous  sulphate 462 

Chlorates  and  nitrates,  determination  of,  iodometric  and  gas-volumetric  273 
Chlorine,  determination  of,  in  alkali  chlorides  and  iodides,  by  distilla- 
tion with  ferric  sulphate 449 

with  nitrite 451 

in  benzol  derivatives 448 

with  bromine,  in  alkali  chlorides  and  iodides  452 

fixation  of,  on  silver  anode 20 

gravimetric  determination  of,  by  liberation  of  iodine  and  ab- 
sorption of  that  element  by  silver 443 

bromine  and  iodine,  detection  of 440 

Chromium,  gravimetric  estimation  of,  as  silver  chromate 406 

separation  of,  from  iron  by  gaseous  hydrogen  chloride 508 

in  chromic  condition,  volumetric  estimation  of 413 

Chromic  acid,  estimation  of,  with  ferric  iron  and  vanadic  acid 510 

iodometric  estimation  of 407 

and  vanadic  acid 409,  41 1 

Chromic  sulphate,  hydrolysis  of,  in  iodide-iodate  mixture 63 

Cobalt,  separation  of,  from  nickel,  by  ether-hydrochloric  acid 492 

Cobalt  sulphate,  hydrolysis  of,  in  iodide-iodate  mixture 63 

Copper,  determination  of,  as  cuprous  iodide 114 

by  titration  of  precipitated  oxalate 123,  125 

electrolytic  determination  of,  on  rotating  cathode 116 

on  rotating  cathode  of  silver 117 

gravimetric  determination  of,  as  sulphocyanate 108 


530  INDEX  OF  SUBJECTS 

PAGE 

Copper,  iodometric  estimation  of 1 18 

iodometric  estimation  of,  associated  with  arsenic 318 

with  antimony 318 

precipitation  as  oxalate 125 

separation  of,  from  cadmium,  by  precipitation  as  cuprous  iodide  1 14 

as  oxalate,  associated  with  lead 135 

from  arsenic,  cadmium,  iron,  tin,  zinc  131 
by  desiccation  process  .  132 
by  acetic  acid  process. .  134 
as  sulphocyanate,  from  antimony,  arsenic,  bis- 
muth, and  tin 112 

Copper  oxalate,  solubility  of 125 

prevention  of  supersaturation  of  solutions  of 129 

Dithionic  acid  and  dithionates,  determination  of 369 

Electrolysis  with  the  filtering  crucible,  with  filtration 16 

continuous 17 

subsequent 13 

Electrolytic  processes 1 1 

Ferricyanides,  detection  of,  with  ferrocyanides  and  sulphocyanides . . .  .275,  276 

Ferric  chloride,  behavior  of,  in  the  Jones  reductor 497 

Ferric  sulphate,  hydrolysis  of,  in  iodide-iodate  mixture 63 

Ferrocyanides,  detection  of,  with  ferricyanides  and  sulphocyanides. .  .275,  276 

Ferrous  sulphate,  hydrolysis  of,  in  iodide-iodate  mixture 63 

titration  of,  in  presence  of  nitric  acid 498 

Filtering  crucible,  in  electrolytic  analysis 13 

Fluorine,  detection  of 432 

iodometric  determination  of 439 

evolved  as  silicon  fluoride,  estimation  of 436 

Fluosilicic  acid,  acidimetric  estimation  of 432 

in  alcoholic  solution 433 

in  water  solution 434 

iodometric  estimation  of 435 

Force  pump  (Kreider) 8 

Gaseous  products,  determination  of,  by  loss I 

prevention  of  mechanical  loss  in  evolution  of 6 

transfer  under  pressure 7 

Gold,  determination  of,  colorimetric 150 

electrolytic 145 

iodometric  (small  amounts) 146 

Halogens,  determination  of,  by  electrolytic  reduction  of  silver  salts.  .  .  .  459 

in  benzol  derivatives 447 

Hydrochloric  acid,  electrolysis  of,  with  silver  anode 20 

Hydrolysis,  in  bromide-bromate  mixture 70 

in  iodide-iodate  mixture 61 

Hydrogen,  determination  of,  by  loss 2 

lodic  acid,  use  of,  in  analysis  of  iodides 454 

Iodides,  determination  of,  by  liberation  of  iodine  and  absorption  of  that 

element  by  silver 446 

by  use  of  iodic  acid 454 

Iodide-iodate  mixture,  reaction  of,  with  salts  —  alums 67 

aluminium  sulphate ...  62 

ammonium  sulphate. . .  65 

chromic  sulphate 63 

cobalt  sulphate 63 

iron  sulphates 63 

nickel  sulphate 64 


INDEX  OF  SUBJECTS  531 

PAGE 

lodide-iodate  mixture,  reaction  of,  with  salts  —  stannic  sulphate 63 

zinc  sulphate 64 

use  of,  in  determination  of  free  acids 59 

of  acids  liberated  by  hydrolysis 61 

of  alkali  hydroxides  and  carbonates ....  60 

Iodine,  gravimetric  determination  of,  by  absorption  in  silver 444 

reaction  of,  with  alkali  hydroxides 70 

standardization  of,  by  arsenic  trioxide 29 

by  silver 27 

in  haloid  salts,  iodometric  determination  of 457 

in  iodides,  gravimetric  determination  of,  by  liberation  and  ab- 
sorption in  silver 446 

bromine  and  chlorine,  detection  of 440 

Iodometric  processes 27 

Iron,  estimation  of,  in  presence  of  titanium 499 

separation  of,  from  copper  precipitated  as  oxalate  ...  .131,  132,  133,  134 

by  gaseous  hydrogen  chloride 504 

from  aluminium 506 

from  beryllium 507 

from  chromium 507 

from  zirconium 508 

(ferric),  estimation  of,  in  presence  of  vanadium 508 

vanadic  acid  and  chromic  acid  510 

reduction  of  titanous  sulphate,  and  estimation  by  per- 
manganate   502 

volumetric  determination,  by  thiosulphate  and  iodine ....  492 
(ferrous;,  permanganate  titration  of,  in  presence  of  chlorides.  .  .  .48,  497 

of  nitric  acid.  .  .  .  498 

analysis,  standardization  of  permanganate  in 495 

chloride,  behavior  of,  in  Jones  reductor 497 

Lanthanum,  estimation  of,  by  permanganate  titration,  precipitated  as 

oxalate 218 

Lead,  detection  of,  separated  as  sulphate 252 

electrolytic  determination  of,  as  dioxide 252 

estimation  of,  by  permanganate  titration  of  oxalate 254 

Liquids,  transfer  of,  under  pressure 7 

Magnesium,  determination  of,  by  precipitation  as  ammonium  magne- 
sium carbonate  and  ignition 154 

as  pyrophosphate 156 

separation  of,  from  alkalies,  by  arsenate  process 158 

Manganese,  determination  of,  as  oxide 478 

as  pyrophosphate 482 

as  sulphate 477 

electrolytic  determination  of ' 485 

precipitation  of,  by  chlorate  process,  and  determination ....  487 

separation  of,  as  carbonate,  and  determination  as  oxide ....  481 

Mechanical  processes I 

Mercury,  determination  of,  by  permanganate  titration 

after  precipi  tation  with  ammonium  oxalate  1 97 

of  mercurous  salts 198 

by  titration  with  thiosulphate  and  iodine.  .  .  196 

gravimetric,  as  mercurous  oxalate 195 

separation  of,  from  aluminium  by  ether-hydrochloric  acid,  and 

determination 217 

Molybdic  acid,  estimation  of,  by  reduction  in  Jones  reductor,  use  of 

ferric  alum,  and  permanganate  titration  of  residue,  424,  426 
gravimetric  estimation  of,  by  liberation  of  iodine  and 

absorption  of  that  element  by  silver 414 


532  INDEX  OF  SUBJECTS 

PAGE 

Molybdic  acid,  iodometric  estimation  of 415 

by  digestion  method 415 

by  distillation  process 416 

by  reduction  with  hydriodic  acid 
and  reoxidation  of  residue,  by 

iodine 420 

by  reduction  with  hydriodic  acid 
and  reoxidation  of  residue,  by 

permanganate 42 1 

and  vanadic  acid,  determination  of,  by  reductions  and 

oxidations 427 

Neutralization  processes,  with  use  of  acid  anhydrides  as  standards 56 

organic  acids  as  standards 56 

succinic  acid,  as  standard 54 

Nickel,  detection  of,  in  presence  of  cobalt 491 

determination  of,  electrolytic 489 

by  precipitation  as  oxalate,  and  permanganate 

titration -.•••: 49° 

separation  of,  from  cobalt,  by  ether-hydrochloric  acid 492 

Nickel  sulphate,  hydrolysis  of,  in  iodide-iodate  mixture 64 

Nitrates,  estimation  of,  by  ignition  with  sodium  paratungstate 256 

by   reduction   with   ferrous  chloride   and   gas- 
volumetric  estimation  of  nitrogen  dioxide ....  260 
by  reduction  with   ferrous  sulphate  and   per- 
manganate titration 258 

iodometric  estimation  of,  by  action  of  manganous  chloride. . . .  263 
by   distillation   with   antimony   tri- 
chloride . 268 

by     distillation     with     iodide     and 

phosphoric  acid 266 

Nitrates  and  chlorates,  iodometric  and  gas- volumetric  estimation  of ....  273 

Nitrates  and  nitrites,  iodometric  and  gas  volumetric  estimation  of 272 

Nitrites,  gas-volumetric    determination    of,    with    use    of    manganous 

chloride • ...'..  271 

iodometric  determination  of,  with  use  of  manganous  chloride.  .  271 

with  potassium  iodide  and  arsenite. .  269 

Nitrites  and  nitrates,  iodometric  and  gas- volumetric  determination  of.  .  272 

Nitrogen,  determination  of,  liberated  by  hypobromite .  256 

Organic  acids  and  anhydrides,   use  of,  as  standards  in  neutralization 

processes 54 

benzoic  acid 58 

malonic  acid 57 

phthaiic  acid 58 

succinic  acid 57 

phthaiic  anhydride 58 

succinic  anhydride 57 

Organic  substances,  combustion  of,  in  the  wet  way 234 

by  chromic  acid 236 

by  permanganate 234 

Oxidation  processes 41 

Oxidizers,  gravimetric  determination  of,  by  liberation  of  iodine    and 

absorption  of  that  element  by  silver 361 

Oxygen,  amount  used  in  oxidation  by  chromic  acid,  and  indirect  estima- 
tion of  oxygen  content  of  organic  substances 239 

iodometric  estimation  of,  in  air 355 

in  water  solution 360 

loss  of,  in  oxidations  by  permanganate 42 


INDEX  OF  SUBJECTS  533 

PAGE 
Perchlorates,  detection   of,    associated    with   chlorides,    chlorates   and 

nitrates 465 

iodometric  determination  of 467 

Perchloric  acid,  preparation  of,  for  potassium  determination 88 

for  sodium  test 76 

Permanganate,  standardization  of 41,  42,  362,  495 

Persulphates,  determination  of,  by  arsenate-iodide  method 370 

by  method  of  Griitzner 372 

by  method  of  LeBlanc  and  Eckardt. ...  371 

by  method  of  Mondolfo 374 

by  method  of  Namias 374 

Phosphoric  acid,  estimation  of,  as  magnesium  pyrophosphate 282 

by  permanganate,  precipitated  as  uranyl 

phosphate 286 

iodometric  estimation  of,  precipitated  as  ammonium 

phosphomolybdate 285 

Phosphorus  in  iron,  iodometric  determination  of,  precipitated  as  ammo- 
nium phosphomolybdate 283 

Potassium,  detection  of,  spectroscopic 80 

effect  of  sodium  salts  on 82 

determination  of,  spectroscopic 83 

in  presence  of  sodium  salts 85 

estimation  of,  as  cobalti-nitrite  by  gravimetric  process.  ...  95 

by  volumetric  process 93,  95 

in  animal  fluids  (blood,  lymph,  milk,  urine) .  .  98 

in  fertilizers 96 

in  mixtures  of  salts 95 

in  pure  salt. 94 

in  soils 97 

separation  of,  and  estimation  as  perchlorate 88 

as  pyrosulphate 92 

Potassium  permanganate,  reaction  of,  as  affected  by  concentration  of 

acid,  time  and  temperature.  . .  .42-48 
upon   ferrous   salt,    with    hydro- 
chloric acid 48,  50 

upon  oxalic  acid,  with  hydrochlo- 
ric acid 50 

with  other  chlo- 
rides    52 

standardization  of,  by  arsenic  trioxide 41 

with  iodine  42 
by  liberation  of  iodine  and 
absorption  of  that  ele- 
ment by  silver 42,  362 

in  iron  analysis 495 

Precipitates,  purification  of,  by  solution  and  reprecipitation I  o 

Rotary  shaker 9 

Rotating  cathode ii 

Rubidium,  estimation  of,  as  acid  sulphate 106 

spectroscopic  determination  of,  in  pure  salt 102 

in  presence  of  potassium  salt  105 

effect  of  potassium  on 104 

effect  of  sodium  on 104 

Selenic  acid,  iodometric  determination  of,  by  action  of  hydrochloric  acid 

by  action  of  hydrobromic  acid  3* 

by  action  of  hydriodic  acid . .  388 


534  INDEX  OF  SUBJECTS 

PACK 

Selenious  acid,  gravimetric  determination  of,  by  precipitation  of  selenium  376 

by  liberation  of  iodine  and 
absorption  of  that  ele- 
ment by  silver 375 

iodometric  estimation  of 377 

by  contact  method 377 

by  differential  method 380 

by  distillation  method 379 

volumetric  estimation  of,  by  permanganate 382 

by  thiosulphate 383 

Selenium,  separation  of,  from  tellurium,  and  estimation,  by  differential 

volatility,  of  bromides 390 

Silicon,  detection  of,  in  silicates  and  fluosilicates 241 

Silicon  fluoride,  estimation  of,  eliminated  at  high  temperature 436 

Silver,  electrolytic  determination  of 138 

gravimetric  determination  of,  as  chromate 136 

iodometric  estimation  of,  by  precipitation  as  chromate 140 

reduced  by  arsenite  from  the  chloride  .  .  143 

anode,  use  of,  for  fixation  of  chlorine 20 

chloride,  electrolytic  reduction  of 460 

bromide,  electrolytic  reduction  of 460 

iodide,  electrolytic  reduction  of 461 

Sodium,  detection  of,  by  hydrochloric  acid  in  alcohol 74 

after  separation  of  potassium  perchlorate 75 

in  mixtures  of  salts  of  other  elements 78 

estimation  of,  as  pyrosulphate 79 

Sodium  chloride,  electrolysis  of,  with  mercury  cathode 22,  26 

electrolytic  analysis  of,  alkalinity  of  inner  cell 23 

transfer  of  silver  to  cathode ...  24 

Stannic  chloride,  hydrolysis  of,  in  iodide-iodate  mixture 63 

Starch  indicator,  colors  of,  with  free  iodine 29,  32 

delicacy  of,  in  presence  of  potassium  iodide 35,  38 

effect  of  temperature  upon 38 

effects  of  varying  amounts  of 36 

end  reaction  of,  with  tartar  emetic 37 

from  various  sources 36 

loss  of  iodine  in  action  of < 30,  36 

preparation  of 33 

Strontium,  detection  of,  associated  with  calcium  and  lead 160 

gravimetric  estimation  of,  precipitated  as  oxalate 180 

volumetric  estimation  of,  precipitated  as  oxalate 182,  183 

Strontium  and  barium,  separation  of,  and  determination,  by  action  of 

amyl  alcohol  on  the  bromides 167 

Strontium  and  calcium,  detection  of,  by  action  of  amyl  alcohol  on  the 

nitrates 163 

separation  of,  and  estimation,  by  action  of  amyl 

alcohol  on  the  nitrates 164 

Succinic  acid,  use  of,  as  standard,  in  iodometric  processes 56 

in  neutralization  processes 54 

Sulphides,  detection  of,  associated  with  sulphites,  sulphates  and  thio- 

sulphates 363 

Sulphites,  detection  of,  associated  with  sulphides,  sulphates  and   thio- 

sulphates 363 

iodometric  determination  of,  in  alkaline  solution 366 

Sulphocyanates,  detection  of,  with  ferricyanides  and  ferrocyanites . . .  .275,  276 

gravimetric  determination  of 276 

volumetric  estimation  of,  by  permanganate 279 

Tartar  emetic,  preparation  of 36,  40 

titration  of,  by  iodine 38 


INDEX  OF  SUBJECTS  535 

PAGE 

Tartar,  emetic  use  of,  as  an  iodometric  standard 40 

variation  in  composition  of 40 

Telluric  acid,  iodometric  estimation  of 401 

Tellurium,  gravimetric  estimation  of,  as  dioxide  precipitated  by  ammonia 

and  acetic  acid 402 

separation  of,  from  selenium 404 

Tellurous  acid,  estimation  of,  by  titrimetric  precipitation,  as  tellurous 

iodide 398 

gravimetric  estimation  of,  by  liberation  of  iodine  and 

absorption  of  that  element  by  silver 394 

iodometric  estimation  of 399 

volumetric  estimation  of,  by  permanganate 394 

in  presence  of  chloride 396 

in  presence  of  bromide 397 

Thallium,  determination  of,  as  acid  sulphate  and  as  neutral  sulphate ....  219 

gravimetric  estimation  of,  as  chromate 221 

precipitated    as    thallic  hydroxide 
by   potassium  ferricyanide    and 

potassium  hydroxide 220 

iodometric  estimation  of,  precipitated  by  potassium  dichro- 

mate 222 

Thiosulphates,  detection  of,  in  association  with  sulphides,  sulphites  and 

sulphates 363 

iodometric  estimation  of,  effects  of  acid,  concentration 

and  temperature 364 

Tin,   detection    of,    associated    with    arsenic,    by    hydrochloric    acid, 

and  iodide 313 

and  bromide .  . 316 

electrolytic  determination  of 251 

estimation  of,  associated  with  arsenic  and  antimony,  by  ferricyanide 

and  permanganate 322,  323 

separation  of,  from  copper  precipitated  as  oxalate 131 

from  copper  precipitated  as  sulphocyanate 112 

Tin  chloride,  hydrolysis  of,  in  iodide-iodate  mixture 63 

Titanium,  determination  of,  by  reduction  and  titration  with  permanga- 
nate   242 

Uranium,  determination  of,  by  Jones  reductor 430 

Valve  (Kreider) 7 

Vanadic  acid,  estimation  of,  gravimetric,  by  action  of  hydriodic  acid  and 

absorption  of  iodine  by  silver 325 

by  precipitation  as  ammonium  vanadate. .  .  326 

by  precipitation  as  silver  meta- vanadate ..  328 

in  association  with  antimonic  acid 350 

with  arsenic  acid 350 

with  chromic  acid 411 

with  iron 508 

with  iron  and  chromic  acid  510 
with    chromium,     iron    and 

molybdenum 352 

by  permanganate  after  reduction 

by  zinc 346 

with  use  of  ferric  sulphate 349 

with  use  of  silver  sulphate 348 

iodometric  estimation  of,  by  action  of  hydrochloric  acid .  .  330 

by  action  of  hydrobromic  acid  335,  345 

by  action  of  hydriodic  acid .  .  .337,  343 

by  reduction  with  organic  acids 

and  reoxidation  by  iodine  ...  341 


536  INDEX  OF  SUBJECTS 

PAGE 
Vanadium  in  tetroxide  condition,  estimation  of,  by  ferricyanide  and 

permanganate 352 

Volatile  products,  distillation  of,  and  absorption 4,  5 

and  condensation 3 

removal  of,  without  mechanical  loss  of  non-volatile 

material 6 

Zinc,  electrolytic  determination  of 186 

estimation  of,  as  pyrophosphate 185 

precipitation  of,  as  oxalate,  and  estimation 187 

separation  of,  from  aluminium 216 

from  copper  precipitated  as  oxalate 131 

Zinc  chloride,  conversion  of,  to  oxide 186 

Zinc  sulphate,  hydrolysis  of,  in  iodide-iodate  mixture 64 

Zirconium,  separation  of,  from  iron  volatilized  in  hydrogen  chloride.  .244,  508 


14  DAY  USE 

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